Atoms, elements and compounds

Specification

AQA: 4.1.1.1

• use the names and symbols of the first 20 elements in the periodic table, the elements in Groups 1 and 7, and other elements in this specificationame compounds of these elements from given formulae or symbol equations

• write word equations for the reactions in this specification

• write formulae and balanced chemical equations for the reactions in this specification.
(HT) write balanced half equations and ionic equations where appropriate.

Edexcel: 1.18 | 1.19 | 1.25 | 1.25 | 1.28

1.18 — understand how elements are arranged in the Periodic Table: • in order of atomic number • in groups and periods.

1.19 — understand how to deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table

1.25 — write word equations and balanced chemical equations (including state symbols): • for reactions studied in this specification • for unfamiliar reactions where suitable information is provided.

1.26 — calculate relative formula masses (including relative molecular masses) (Mr) from relative atomic masses (Ar)

1.28 — understand how to carry out calculations involving amount of substance, relative atomic mass (Ar) and relative formula mass (Mr)

Cambridge: 2.2.6 | 3.2.2 | 3.3.3 (E) | 8.1.1 | 8.1.4

2.2.6 — State that: (a) Group VIII noble gases have a full outer electron shell (b) the number of outer shell electrons is equal to the group number in Groups I to VII (c) the number of occupied electron shells is equal to the period number

3.2.2 — Define relative molecular mass, Mr, as the sum of the relative atomic masses. Relative formula mass, Mr, will be used for ionic compounds

3.3.3 (E) — Use the relationship mass (g) amount of substance (mol) = molar mass (g / mol) to calculate: (a) amount of substance (b) mass (c) molar mass (d) relative atomic mass or relative molecular / formula mass (e) number of particles, using the value of the Avogadro constant

8.1.1 — Describe the Periodic Table as an arrangement of elements in periods and groups and in order of increasing proton number / atomic number

8.1.4 — Explain similarities in the chemical properties of elements in the same group of the Periodic Table in terms of their electronic configuration

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Atoms, Elements & Isotopes: Cracking the Periodic Table

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

Structure of an Atom

• An atom consists of a central nucleus containing protons and neutrons, with electrons orbiting the nucleus.
• Hydrogen is the simplest element, with just one proton in its nucleus and one electron orbiting it.
• Helium is the next smallest element, with two protons and two neutrons in the nucleus and two electrons orbiting it.

The Periodic Table

• There are around 100 different elements, all organised into the periodic table where each box represents a different element.
• Each element is represented by a nuclear symbol, which includes a one- or two-letter symbol, the atomic number, and the mass number.
• Some element symbols are not obvious from their names — for example, gold is Au and iron is Fe.

Atomic Number

• The atomic number is the number of protons in the nucleus of an atom and is unique to each element.
• Any atom with a given number of protons must be the element that corresponds to that atomic number — for example, any atom with six protons is carbon.
• If an atom has three protons, the periodic table tells us it must be lithium, which has an atomic number of three.

What Are Isotopes?

• Isotopes are different forms of the same element that have the same number of protons but a different number of neutrons.
• The number of neutrons in an atom can vary without changing which element it is, as long as the number of protons stays the same.
• Because isotopes only diﬀer in the number of neutrons, they behave in almost identical ways chemically.

Isotopes of Carbon

• Carbon-12, the most common form of carbon, has six protons, six neutrons, and six electrons, giving it a mass number of 12.
• Carbon-13 is a rarer isotope that still has six protons and six electrons, but has seven neutrons, giving it a mass number of 13.
• The mass number of an isotope is the total number of protons and neutrons in its nucleus.

Isotopic Abundance

• Isotopic abundance refers to how common or rare a particular isotope is within a sample of an element, usually expressed as a percentage.
• Copper has two stable isotopes: copper-63 with an abundance of 69.2% and copper-65 with an abundance of 30.8%.
• The abundances of all isotopes of an element must add up to 100%, as they account for all atoms of that element.

Calculating Relative Atomic Mass

• The relative atomic mass (Aᵣ) is the average mass of all the isotopes of an element, taking into account their relative abundances.
• The formula for relative atomic mass is: Aᵣ = (Σ(isotope abundance × isotope mass) / Σ(abundances of all isotopes))
• For copper, the calculation is ((69.2 × 63) + (30.8 × 65) / 69.2 + 30.8) = (6361.6 / 100) = 63.6 (to one decimal place).
• The relative atomic mass of copper is 63.6, meaning the average mass of a copper atom is 63.6 atomic mass units.

Exam Focus – RAM

Relative Atomic Mass: From Definition to Calculation

Supporting Material

• Atoms are incredibly small and light — for example, a proton or neutron weighs approximately 1.7 imes 10⁻²⁷ kg — making direct mass comparisons impractical.

• It is more useful to compare the mass of an atom to a standard atom, giving us the concept of relative atomic mass (Ar).

• Electrons contribute very little to atomic mass, so the majority of an atom’s mass comes from protons and neutrons in the nucleus.

• Relative atomic mass (Ar) is defined as the average mass of one atom of an element relative to ¹⁄₁₂ the mass of one atom of carbon-12.

•Carbon-12 is chosen as the standard atom because it is the most commonly occurring isotope of carbon and provides a universal reference for chemists worldwide.

• Carbon-12 has the nuclear symbol ¹²₆C, meaning it has 6 protons and 6 neutrons, giving a mass number of 12.

• ¹⁄₁₂ the mass of a carbon-12 atom is approximately equal to the mass of one proton or one neutron, since protons and neutrons have essentially the same mass.

• Because different isotopes of an element exist in nature, the relative atomic mass is the weighted average of all naturally occurring isotopes.

• If an element has only one isotope, its Ar value is equal to its mass number, since there is no variation to average out.

• Chlorine (Cl) exists as two naturally occurring isotopes: ³⁵Cl (75% abundance) and ³⁷Cl (25% abundance). Because chlorine has two isotopes, its Ar value is not a whole number but a weighted average, giving Ar (Cl) = 35.5.

• Isotope abundance refers to the relative proportion of each isotope that is naturally present in a sample of that element.

Exam Question

1 (a) Atoms are made of protons, neutrons and electrons. Atoms of the same element are known as isotopes.

Magnesium-24:

24 12 Mg

Magnesium-25:

25 12 Mg

Complete the table below to show the numbers of electrons, neutrons and protons in these isotopes of magnesium.

(2 marks)

isotope	number of electrons	number of neutrons	number of protons
Magnesium-24
Magnesium-25

2 (b) Boron, B, has two isotopes.

(i) State the meaning of the term isotopes.

(2 marks)

(ii) Table 2.1 shows the relative masses and the percentage abundances of the two isotopes of boron.

Calculate the relative atomic mass of boron to one decimal place.

relative atomic mass = …………………………

(2 marks)

relative mass of isotope	percentage abundance of isotope
10	20
11	80

Table 2.1

(iii) Lithium, sodium and potassium are Group I elements.

Lithium has two naturally occurring types of atoms,

Boron-10:

10 B

Boron-11:

11 B

Complete Table 2.2 to show the numbers of protons, neutrons and electrons in an atom of Boron-11.

(2 marks)

number of protons	number of neutrons	number of electrons

Table 2.2

Determine the relative abundance of Boron-10 present in boron. Give your answer as a percentage.

………………………… %

(1 marks)

Recall one or more pieces of information.

Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons, and therefore different mass numbers.

1 (a) Atoms are made of protons, neutrons and electrons. Atoms of the same element are known as isotopes.

Magnesium-24:

24 12 Mg

Magnesium-25:

25 12 Mg

Complete the table opposite to show the numbers of electrons, neutrons and protons in these isotopes of magnesium.

(2 marks)

isotope	number of electrons	number of neutrons	number of protons	1 mark per correct row
Magnesium-24	12	12	12	✔
Magnesium-25	12	13	12	✔

2 (b) Boron, B, has two isotopes.

(i) State the meaning of the term isotopes.

different atoms of the same element with the same number of protons
different numbers of neutrons

(2 marks)

(ii) Table 2.1 shows the relative masses and the percentage abundances of the two isotopes of boron.

Calculate the relative atomic mass of boron to one decimal place.

(10×20) + (11 x80) =1080
(1080 ÷ 100) = 10.8

(2 marks)

relative mass of isotope	percentage abundance of isotope
10	20
11	80

Table 2.1

(iii) Boron and aluminium are Group III elements.

Boron has two naturally occurring types of atoms,

Boron-10:

10 B

Boron-11:

11 B

Complete Table 2.2 to show the numbers of protons, neutrons and electrons in an atom of Boron-11.

number of protons	number of neutrons	number of electrons
5	6	5

Table 2.2

5p and 5e
6n

(2 marks)

Determine the relative abundance of Boron-10 present in boron. Give your answer as a percentage.

20 %

(1 marks)

Check Point

C________ has six protons, six neutrons, and six electrons

Carbon-12 Carbon 12

Carbon-13 is an isotope of carbon that has six protons and s____ neutrons

seven 7

Different isotopes of the same element have different masses because they contain different numbers of n_______

neutrons

H_____ has two protons and two neutrons in its nucleus, with two electrons orbiting it

Helium

H_______ is the smallest element, with one proton and one electron

hydrogen

I_______ are different forms of the same element that have the same number of protons but a different number of neutrons

isotopes

The number of p______ in an atom determines which element it is

protons

The relative atomic mass is calculated using the equation: sum of (isotope abundance x isotope mass) d______ by the sum of all isotope abundances

divided

The relative atomic mass of copper, calculated from its two stable isotopes, is 6_

There are around 1__ different elements, organised into the periodic table

100 one hundred

Mixtures

Specification

AQA: 4.1.1.2

• describe, explain and give examples of the specified processes of separation

• suggest suitable separation and purification techniques for mixtures when given appropriate information.

Edexcel: 1.10 | 1.25 | 1.31

1.10 — describe these experimental techniques for the separation of mixtures: • simple distillation • fractional distillation • filtration • crystallisation • paper chromatography.

1.31 — understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Cambridge: 12.4.1 | 12.4.2

12.4.1 — Describe and explain methods of separation and purification using: (a) a suitable solvent (b) filtration (c) crystallisation (d) simple distillation (e) fractional distillation

12.4.2 — Suggest suitable separation and purification techniques, given information about the substances involved

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Molecules, Compounds & Mixtures: How Atoms Exist in the Real World

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

What is a Molecule?

• A molecule is a group of two or more atoms held together by chemical bonds.
• Molecules can consist of atoms of the same element, such as oxygen (O₂), or different elements, such as water (H₂O).
• A single atom, like helium, does not count as a molecule because it requires at least two atoms.

What is a Compound?

• A compound is a substance that contains two or more different elements chemically bonded together.
• Water (H₂O) and carbon dioxide (CO₂) are compounds because they each contain more than one type of element.
• Molecules such as oxygen (O₂), chlorine (Cl₂), and nitrogen (N₂) are not compounds because they only contain one type of element.
• In a compound, the elements are always present in fixed proportions, for example water always has two hydrogen atoms for every one oxygen atom.

Chemical Formulae

• Chemical formulae use element symbols and subscript numbers to show how many atoms of each element are in a molecule, e.g.
• H₂O means two hydrogen atoms and one oxygen atom.
• If only one atom of an element is present, no subscript number is written, as seen with the single carbon in CO₂.
• A more complex example is H₂SO₄ (sulfuric acid), which contains two hydrogen atoms, one sulfur atom, and four oxygen atoms.
• Brackets in formulae group atoms together, and the number outside the bracket shows how many of those groups are present, e.g.
• Ca(OH)₂ means one calcium and two groups of OH.

Molecular vs Non-Molecular Compounds

• Some compounds, such as sodium chloride (NaCl), exist as giant structures containing millions or billions of atoms rather than small individual molecules.
• Compounds that form large structures rather than molecules typically contain ionic bonds between their atoms.
• The formula of a non-molecular compound acts as a ratio; for example, NaCl shows that sodium and chloride ions are present in a 1:1 ratio.

What is a Mixture?

• A mixture consists of two or more substances that are not chemically combined together.
• In a mixture, the individual substances are physically combined but no chemical bonds form between them.
• Because the substances in a mixture are not chemically bonded, they retain their individual properties and can be separated by physical methods.

Separating Mixtures Filtration

Separating Mixtures crystallisation

Separating Mixtures distillation

Separating Mixtures Chromatography

Separating Mixtures Chromatogram

• Mixtures can be separated using physical methods such as filtration, crystallisation, distillation or Chromatography.
• These separation techniques work because the components of a mixture are not chemically bonded and have different physical properties.

Key Differences: Molecules, Compounds & Mixtures

• A molecule contains two or more atoms bonded together, whereas a compound specifically requires two or more different elements.
• All compounds are chemically bonded substances with fixed ratios of elements, while mixtures have no chemical bonds and can vary in composition.
• Compounds cannot be separated by physical methods, whereas the components of a mixture can be separated physically.

Check Point

A c_______ is a substance that contains two or more different elements held together by chemical bonds

compound

In a chemical formula, a s________ number indicates how many atoms of that element are present in one molecule

subscript

In chemical formulas containing brackets, the number o______ the brackets indicates how many of those groups are present

outside

In the formula H____, each molecule of sulfuric acid contains two hydrogen atoms, one sulfur atom, and four oxygen atoms

H₂SO₄

Oxygen exists as d_______ molecules, meaning two oxygen atoms are bonded together

diatomic

Oxygen, chlorine, and nitrogen are molecules but not compounds because they contain only one type of e______

element

The chemical formula for carbon dioxide is C__, indicating one carbon atom and two oxygen atoms

CO₂

Sodium chloride is a non-molecular compound that exists as a giant i____ structure containing millions of atoms

ionic

The chemical formula for water is H__, indicating two hydrogen atoms and one oxygen atom

H₂O

The substances in a mixture can be separated using physical methods such as filtration, crystallisation, or d___________

distillation

The development of the model of the atom

Specification

AQA: 4.1.1.3

• describe why the new evidence from the scattering experiment led to a change in the atomic model

• describe the difference between the plum pudding model of the atom and the nuclear model of the atom.

Edexcel: 1.16

1.16 — know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

Cambridge: 2.2.1 | 2.2.2 | 2.2.3 | 2.2.4 | 2.2.5

2.2.1 — Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells

2.2.2 — State the relative charges and relative masses of a proton, a neutron and an electron

2.2.3 — Define proton number / atomic number as the number of protons in the nucleus of an atom

2.2.4 — Define mass number / nucleon number as the total number of protons and neutrons in the nucleus of an atom

2.2.5 — Determine the electronic configuration of elements and their ions with proton number 1 to 20, e.g. 2,8,3

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The History of the Atom: From Ancient Greece to Modern Theory

🎯 Knowledge and Comprehension

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🎯🎯🎯 Application and Analysis

Democritus and Early Atomic Theory

• The ancient Greek philosopher Democritus (around 500 BC) first proposed that all matter is made up of tiny, indivisible particles called atoms, separated by empty space.
• This idea remained largely unchanged for approximately 2,300 years until scientists in the 1800s began to develop and improve upon it.

Dalton’s Solid Sphere Model

• In the early 1800s, John Dalton described atoms as solid spheres, proposing that different types of spheres corresponded to the different chemical elements.
• Dalton’s model was a significant step forward as it linked the concept of atoms directly to the elements we observe in chemistry.

Thomson’s Plum Pudding Model (1897)

• In 1897, J.J. Thomson conducted experiments that proved atoms could not be solid spheres, leading to the discovery of negatively charged particles called electrons.
• Thomson proposed the plum pudding model, in which the atom was described as a ball of diffuse positive charge with discrete electrons embedded within it.

Rutherford’s Gold Foil Experiment (1909)

• Ernest Rutherford and his students fired positively charged alpha particles at a thin sheet of gold foil, expecting them to pass straight through if Thomson’s model were correct.
• Surprisingly, some alpha particles were deflected sideways and a small number bounced directly back, disproving the plum pudding model.
• Rutherford proposed the nuclear model, suggesting that all positive charge is concentrated in a small, dense nucleus at the centre of the atom, surrounded by a cloud of negative electrons.

The Flaw in Rutherford’s Nuclear Model

• A key problem with Rutherford’s nuclear model was that it could not explain why the negatively charged electrons were not pulled into the positively charged nucleus, which would cause the atom to collapse.
• This unresolved issue highlighted the need for a further development in atomic theory.

Bohr’s Atomic Model (1913)

• In 1913, Niels Bohr proposed that electrons orbit the nucleus in fixed energy levels called shells, in a similar way to how planets orbit the Sun.
• The orbiting motion of electrons in shells explained why they do not spiral into the nucleus, solving the major flaw in Rutherford’s model.
• Bohr’s model has been supported by many subsequent experiments and forms the basis of the atomic model used in GCSE Chemistry today.

Discovery of Protons and Neutrons

• Further experiments by Rutherford revealed that the positive charge within the nucleus is made up of small, discrete particles called protons.
• James Chadwick later provided experimental evidence for the existence of neutral particles within the nucleus, which are called neutrons.
• Together, protons and neutrons make up the nucleus of an atom, completing our modern understanding of atomic structure.

Check Point

A flaw in Rutherford’s nuclear model was that it could not explain why electrons did not c_______ into the positive nucleus

collapse

Dalton proposed that different types of solid spheres make up the different e_______

elements

Electrons are negatively charged particles that o_____ shells around the nucleus of an atom

occupy

Ernest Rutherford fired positively charged a____ particles at a thin sheet of gold foil in 1909

alpha

In Rutherford’s gold foil experiment, some alpha particles were d________ to the side and a small number were deflected back the way they had come

deflected

John Dalton described atoms as s____ spheres in the 1800s

solid

Niels Bohr proposed in 1913 that electrons orbit the nucleus in f____ shells, similar to how planets orbit the Sun

fixed

The idea that matter is composed of tiny indivisible particles is known as a_____ theory

atomic

The n______ of an atom contains both protons and neutrons

nucleus

Thomson’s experiments showed that atoms contained n_________ charged particles, which we now call electrons

negatively

Relative electrical charges of subatomic particles

Specification

AQA: 4.1.1.4

• Students should be able to use the nuclear model to describe atoms.

Edexcel: 1.16

1.16 — know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

Cambridge: 2.2.1 | 2.2.2 | 2.2.3 | 2.2.4 | 2.2.5

2.2.1 — Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells

2.2.2 — State the relative charges and relative masses of a proton, a neutron and an electron

2.2.3 — Define proton number / atomic number as the number of protons in the nucleus of an atom

2.2.4 — Define mass number / nucleon number as the total number of protons and neutrons in the nucleus of an atom

2.2.5 — Determine the electronic configuration of elements and their ions with proton number 1 to 20, e.g. 2,8,3

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Atoms: Structure, Particles & the Periodic Table

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

What is an Atom?

• Everything — living or non-living — is made up of incredibly tiny particles called atoms.
• Atoms are so small that a single cell contains more than 100 trillion of them.
• Atoms have a radius of around 0.1 nanometres, making them far too small to see with the naked eye.

Basic Structure of an Atom

• An atom consists of a central nucleus surrounded by electrons that orbit in regions called shells.
• The nucleus is made up of two types of particles: protons and neutrons, packed tightly together.
• Most of an atom is empty space, as the distance is measured from the nucleus all the way to the outermost electron shell.
• The nucleus is approximately 10,000 times smaller than the overall width of the atom.

Relative Mass of Subatomic Particles

• Protons and neutrons both have a relative mass of 1, as they are similar in size.
• Electrons are approximately 2,000 times smaller in mass than protons or neutrons, so their relative mass is considered negligible (effectively 0).

Relative Charge of Subatomic Particles

• Protons carry a relative charge of +1 — both ‘proton’ and ‘positive’ begin with the letter P, which helps you remember this.
• Neutrons carry no charge and are electrically neutral — ‘neutron’ and ‘neutral’ share the same beginning.
• Electrons carry a relative charge of −1, equal in magnitude but opposite in sign to the charge of a proton.
• Because protons and neutrons are in the nucleus and electrons are outside it, the nucleus is positively charged overall.

Overall Charge of an Atom

• In a neutral atom, the number of protons always equals the number of electrons, so the positive and negative charges cancel out.
• Because the charges balance, an atom as a whole has no overall electrical charge.

Ions — Charged Atoms

• When an atom gains or loses electrons, the charges no longer balance and the atom becomes a charged particle called an ion.
• If an atom gains electrons, there are more negative electrons than positive protons, forming a negative ion (e.g. a 1− or 2− ion).
• If an atom loses electrons, there are more positive protons than negative electrons, forming a positive ion (e.g. a 1+ ion).

Elements and the Periodic Table

• Each different type of atom is called an element, and every element has its own box in the periodic table.
• Each box in the periodic table is known as a nuclear symbol and contains key information about that element.
• The elemental symbol is a one- or two-letter abbreviation representing the element, such as O for oxygen and Li for lithium.

Atomic Number

• The atomic number, found in the bottom left of the nuclear symbol, tells you the number of protons in an atom of that element.
• It is the number of protons that determines which element an atom is — no two elements have the same atomic number.
• Because a neutral atom has equal numbers of protons and electrons, the atomic number also tells you the number of electrons.

Mass Number and Calculating Neutrons

• The mass number, found in the top left of the nuclear symbol, gives the total number of protons and neutrons in the nucleus.
• The number of neutrons can be calculated using the formula: neutrons = mass number – atomic number.
• For example, oxygen has a mass number of 16 and an atomic number of 8, so it has 16 – 8 = 8 neutrons.
• The number of neutrons does not always equal the number of protons — for instance, lithium has 3 protons but 4 neutrons.

Quick Reference: Subatomic Particles Summary

• Protons: located in the nucleus, relative mass = 1, relative charge = +1.
• Neutrons: located in the nucleus, relative mass = 1, relative charge = 0.
• Electrons: orbit the nucleus in shells, relative mass ≈ 0, relative charge = −1.

Check Point

A positive ion is formed when an atom l____ one or more electrons, giving it more protons than electrons

loses

An atom consists of a central n______ surrounded by electrons that orbit in shells

nucleus

Electrons have a n_________ relative mass and a relative charge of minus one

negligible

Most of the volume of an atom is e____ space between the nucleus and the orbiting electrons

empty

The a_____ number of an element tells us the number of protons in the nucleus of its atoms

atomic

The m___ number of an atom is the total number of protons and neutrons in its nucleus

mass

The nucleus of an atom is made up of protons and n_______ packed tightly together

neutrons

The number of neutrons in an atom is calculated by s__________ the atomic number from the mass number

subtracting

The relative mass of an electron is approximately 2____ times smaller than that of a proton or neutron

2,000 2000, two thousand

When an atom loses or gains electrons, it becomes a charged particle called an i__

ion

Size and mass of atoms

Specification

AQA: 4.1.1.5

• Students should be able to calculate the numbers of protons, neutrons and electrons in an atom or ion, given its atomic number and mass number.

• Students should be able to relate size and scale of atoms to objects in the physical world.

Edexcel: 1.16

1.16 — know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

Cambridge: 2.2.1 | 2.2.2 | 2.2.3 | 2.2.4 | 2.2.5

2.2.1 — Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells

2.2.2 — State the relative charges and relative masses of a proton, a neutron and an electron

2.2.3 — Define proton number / atomic number as the number of protons in the nucleus of an atom

2.2.4 — Define mass number / nucleon number as the total number of protons and neutrons in the nucleus of an atom

2.2.5 — Determine the electronic configuration of elements and their ions with proton number 1 to 20, e.g. 2,8,3

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Ions: How and Why Atoms Gain or Lose Electrons

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

What is an Ion?

• An ion is a charged particle that can be a single atom, such as Na⁺, or a group of atoms, such as the hydroxide ion (^–OH ).
• Ions are formed when atoms gain or lose electrons in order to achieve a full outer shell and become more stable.

The Periodic Table and Outer Shell Electrons

• The group number in the periodic table tells you how many electrons an atom has in its outermost shell — for example, Group 1 elements have one outer electron.
• Atoms need to lose or gain electrons until their outermost shell is full to reach a stable electronic configuration.
• Losing or gaining electrons requires energy, so atoms that only need to move one or two electrons are much more likely to form ions.

Which Groups Commonly Form Ions?

• Group 1 elements (alkali metals) lose one electron to form 1+ ions, as this requires very little energy.
• Group 2 elements lose two electrons to form 2+ ions.
• Group 6 elements gain two electrons to form 2− ions, and Group 7 elements gain one electron to form 1− ions.
• Elements in Groups 3, 4, and 5 would need to lose or gain three or four electrons, which requires too much energy, so they rarely form simple ions.

Writing Equations for Ion Formation

• When an atom loses electrons to form a positive ion, the electron(s) are written on the right-hand side of the equation, e.g.
Na → Na⁺ + e⁻.
• When an atom gains electrons to form a negative ion, the electron(s) are written on the lefthand side, e.g.
Cl + e⁻ → Cl⁻.
• For atoms that lose or gain multiple electrons, the number of electrons is shown accordingly, e.g.
Mg → Mg²⁺ + 2e⁻ for magnesium and O + 2e⁻ → O²⁻ for oxygen.

Check Point

Alkali metals l___ one electron to form 1+ ions

lose

An ion is a c______ particle that can be a single atom or a group of atoms

charged

Elements in groups three, four, and five r_____ form ions because gaining or losing three or four electrons requires too much energy

rarely

Forming ions requires energy, so atoms that only need to lose or gain one or two electrons form ions more r______

readily

Group one elements are called the a_____ metals

alkali

Group s__ elements gain two electrons to form 2- ions

six

Ions are formed when atoms gain or lose e________

electrons

The g____ number in the periodic table indicates how many electrons an atom has in its outermost shell

group

When a chlorine atom gains one electron, it forms a c_______ ion with a charge of 1-

chloride

When a s_____ atom loses an electron, it forms a sodium ion with a charge of 1+

sodium

Relative atomic mass

Specification

AQA: 4.1.1.6

• Students should be able to calculate the relative atomic mass of an element given the percentage abundance of its isotopes.

Edexcel: 1.16 | 1.17

1.16 — know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

1.17 — be able to calculate the relative atomic mass of an element (Ar) from isotopic abundances

Cambridge: 2.3.4 (E) | 3.2.2

2.3.4 (E) — Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

3.2.2 — Define relative molecular mass, Mr, as the sum of the relative atomic masses. Relative formula mass, Mr, will be used for ionic compounds

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Relative Formula Mass & Percentage Mass Calculations

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

Introduction

• The mass number of an element is the total number of protons and neutrons in its nucleus, found in the top left of its nuclear symbol.
• Relative atomic mass (Ar) is the average mass of all the isotopes of an element, taking into account how common each isotope is.
• For example, chlorine has isotopes with mass numbers of 35 and 37, giving it a relative atomic mass of 35.5.

What is Relative Formula Mass

• Relative formula mass (Mr) is the average mass of an entire compound, calculated by adding together the Ar values of all atoms in its molecular formula.
• Mass is always written as Mr and has no units, as it is a relative measure.
• To find Mr , you must account for every atom present in the formula, including any that are multiplied by subscript numbers.

Calculating Mr : Magnesium Chloride Example

• Magnesium chloride has the formula MgCl₂, containing one magnesium atom (Ar = 24) and two chlorine atoms (Ar = 35.5 each).
• The calculation is 24 + (35.5 × 2) = 95, so the Mr of magnesium chloride is 95.

Calculating Mr : Sulfuric Acid Example

• Sulfuric acid has the formula H₂SO₄, containing two hydrogen atoms (Ar = 1), one sulfur atom (Ar = 32), and four oxygen atoms (Ar = 16).
• The calculation is (2 × 1) + (1 × 32) + (4 × 16) = 98, so the Mr of sulfuric acid is 98.

Calculating Percentage Mass of an Element

• The percentage mass of an element in a compound is calculated using the formula: (Ar × number of atoms of that element) / Mr × 100 For sulfur in sulfuric acid: 32 × 198 × 100 = 32.7%, meaning sulfur makes up 32.7% of the mass of sulfuric acid.

Percentage Mass: Calcium Hydroxide Example

• Calcium hydroxide has the formula Ca(OH)₂, meaning the subscript 2 applies to both the oxygen and hydrogen inside the brackets.
• The Ar values are: calcium = 40, oxygen = 16, hydrogen = 1, giving an Mr of 40 + (2 × 16) + (2 × 1) = 74.
• The percentage mass of oxygen in calcium hydroxide is 16 × 274 × 100 = 43.2%.

Key Tips for Exam Success

• Always use the periodic table to look up Ar values rather than trying to memorise them.
• Pay close attention to brackets in chemical formulae, as the number outside the bracket multiplies all atoms inside it.
• Double-check your Mr calculation before using it in a percentage mass question, as an error there will affect your final answer.

Exam Focus – RFM

Relative Formula Mass: From Atoms to Equations

Supporting Material

• Each element’s block on the periodic table shows two numbers: the larger number is the relative atomic mass (Ar) and the smaller number is the atomic number.

• Carbon has an Ar of 12, oxygen has an Ar of 16, and neon has an Ar of 20, as shown on the periodic table.

• Always use the periodic table to look up relative atomic masses rather than trying to memorise them all.

• The relative formula mass (Mr) of a compound is the sum of all the relative atomic masses (Ar) of the atoms present in its chemical formula.

• For small covalent molecules, the term relative molecular mass is sometimes used instead of relative formula mass, but both terms mean exactly the same thing.

• Mr has no units, as it is a relative measure compared to one-twelfth the mass of a carbon-12 atom.

Step 1
Write out the correct chemical formula of the compound you are investigating.
Step 2
Count how many atoms of each element are present in the formula, taking care with brackets and subscript numbers.
Step 3
Look up the Ar of each element from the periodic table.
Step 4
Multiply each Ar by the number of atoms of that element and add all the values together to find Mr.

In the reaction HCl + NaOH → NaCl + H₂O, the Mr values are:

HCl = 36.5

NaOH = 40

NaCl = 58.5

H₂O = 18

The sum of the Mr values for the reactants is 36.5 + 40 = 76.5 and for the products is 58.5 + 18 = 76.5, confirming the equation is balanced.

Exam Question

1 Methylbutanoic acid has the molecular formula C₅H₁₀O₂.

Complete Table 1.1 to calculate the relative molecular mass of C₅H₁₀O₂.

atom	number of atoms	relative atomic mass	Calculation
carbon	5	12	5 × 12 = 60
hydrogen		1
oxygen		16

Table 1.1

relative molecular mass = …………………………

Students should use numbers given in the question to work out the answer.

The reaction between an acid and a base to form a salt plus water.
Word Equation: Acid + Base → Salt + Water
Ionic Equation: H⁺_(aq) + OH^–_(aq)→ H₂O_(l)

1 Methylbutanoic acid has the molecular formula C₅H₁₀O₂.

Complete Table 1.1 to calculate the relative molecular mass of C₅H₁₀O₂.

atom	number of atoms	relative atomic mass	Calculation
carbon	5	12	5 × 12 = 60
hydrogen	10	1	10 × 1 = 10
oxygen	2	16	2 × 16 = 32

Table 1.1

relative molecular mass = 102

102 for 2 marks
if 2 marks not scored 1 mark for 10 (for H)
OR 32 (for 0)

Check Point

Calcium hydroxide has the molecular formula C______

Ca(OH)₂

Chlorine has a relative atomic mass of 35.5 because it is the w_______ average of its isotopes with mass numbers 35 and 37

weighted

Sulfuric acid has the molecular formula H____

H₂SO₄

The mass number of an element is the total number of protons plus n_______ in its nucleus

neutrons

The molecular formula of magnesium chloride is M____

MgCl₂

The percentage mass of oxygen in calcium hydroxide (Ca(OH)2) is 4____, calculated as (16 x 2 / 74) x 100

The percentage mass of sulfur in sulfuric acid (H2SO4) is 3____, calculated as (32 x 1 / 98) x 100

The relative atomic mass (Ar) is the average mass of all the isotopes of an element, weighted by their a________

abundance

The relative atomic mass differs from the mass number because it accounts for the existence of i_______ of that element

isotopes

The relative formula mass of magnesium chloride (MgCl2) is 9_, calculated as 24 + (35.5 x 2)

Electronic structure

Specification

AQA: 4.1.1.7

• represent the electronic structures of the first twenty elements of the periodic table using numbers and diagrams.

Edexcel: 1.16

1.16 — know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

Cambridge: 2.2.1 | 2.2.2 | 2.2.3 | 2.2.5 | 2.2.6

2.2.1 — Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells

2.2.2 — State the relative charges and relative masses of a proton, a neutron and an electron

2.2.3 — Define proton number / atomic number as the number of protons in the nucleus of an atom

2.2.5 — Determine the electronic configuration of elements and their ions with proton number 1 to 20, e.g. 2,8,3

Print Cornell Revision Notes ⧉

Electron Arrangements in Atoms: Shells, Stability & Ions

🎯 Knowledge and Comprehension

🎯🎯 Comprehension and Application

🎯🎯🎯 Application and Analysis

Review

• Every shell is a different energy level.
• The atomic number of an element tells us both the number of protons and the number of electrons in a neutral atom.
• Shells are drawn as rings around the nucleus, and electrons are represented as dots or crosses on these rings.

Rules for Filling Electron Shells

• The first shell, closest to the nucleus and with the lowest energy level, fills first and can hold a maximum of 2 electrons.
• The second and third shells can each hold a maximum of 8 electrons.
• When filling a shell, place one electron in the top, right, bottom, and left positions before adding a second electron to each position (filling clockwise).
• Electrons always fill the innermost available shell before moving to the next one.

Writing Electron Configurations

• The electron arrangement of an atom can be written as a series of numbers separated by commas, each number representing the electrons in successive shells.
• For example, sodium (atomic number 11) has the electron configuration 2, 8, 1, meaning 2 electrons in the first shell, 8 in the second, and 1 in the third.
• You need to be able to work out the electron arrangement of any of the first 20 elements in the periodic table for your GCSE exam.

Stability and the Full Outer Shell

• An atom is considered stable when its outermost shell (the shell furthest from the nucleus) is completely full.
• Most single atoms have incomplete outer shells and are therefore unstable, meaning they will react with other atoms to achieve a full outer shell.
• Atoms react to form molecules or compounds in order to achieve a stable, full outer shell of electrons.

Noble Gases: The Stable Exception

• The noble gases, found in Group 0 of the periodic table, are the exception because they already have completely full outer shells.
• Because noble gases have full outer shells, they have no tendency to gain or lose electrons and so they do not readily react with other elements.
• Neon is an example of a noble gas with a full outer shell, giving it its characteristic chemical inertness.

Worked Example: Argon

• Argon has an atomic number of 18, so it has 18 electrons arranged as 2, 8, 8 across three shells.
• Since 2 + 8 + 8 = 18, argon’s third shell is completely full, making it a stable noble gas.
• This electron arrangement can be written numerically as 2, 8, 8 rather than drawing out all the shells.

Worked Example: Calcium

• Calcium has an atomic number of 20, giving it the electron arrangement 2, 8, 8, 2 across four shells.
• Calcium is unstable because its outer (fourth) shell contains only 2 electrons instead of a full 8.
• If calcium loses its 2 outer electrons, the third shell becomes the outermost shell and is completely full, making the atom stable.
• After losing 2 electrons, calcium still has 20 protons but only 18 electrons, giving it an overall charge of 2+, making it a calcium ion written as Ca²⁺.

Ions and Ionic Charge

• An ion is formed when an atom gains or loses electrons, resulting in an overall positive or negative charge.
• Ions are represented with square brackets around the electron diagram and the overall charge written in the top right corner.
• A positive ion (cation) is formed when an atom loses electrons, whilst a negative ion (anion) is formed when an atom gains electrons.

Worked Example: Fluorine

• Fluorine has an atomic number of 9, giving it the electron arrangement 2, 7, meaning its outer shell has 7 out of a possible 8 electrons.
• To become stable, fluorine needs to gain 1 electron from another atom, giving it a full outer shell of 8 electrons.
• After gaining 1 electron, fluorine has 9 protons but 10 electrons, giving it an overall charge of 1−, making it the fluoride ion F with the electron arrangement 2, 8.

Check Point

An atom is described as u_______ when its outermost shell is not completely full

unstable

Argon has an atomic number of 1_ and an electron arrangement of 2, 8, 8

Atoms with i_________ outer shells will react with other atoms to gain or lose electrons

incomplete

Because n____ gases have full outer shells, they do not react with other elements

noble

Electron arrangements can be represented n__________, for example sodium is written as 2, 8, 1

MgCl₂

Fluorine has an atomic number of _ and an electron arrangement of 2, 7

Sodium has an atomic number of 1_, giving it an electron arrangement of 2, 8, 1

The atomic number of an element tells you both the number of protons and the number of electrons in a n______ atom

neutral

The first electron shell can hold a maximum of t__ electrons

two

When fluorine gains one electron, it forms a f_______ ion with a charge of 1-

fluoride