OCR A Level Chemistry A (H432)

2.1.1 Atomic Structure and Isotopes

1. Specification Coverage

SPEC: 2.1.1 — Atomic structure and isotopes

OCR divides this section into five explicit learning outcomes:

• 2.1.1(a) Isotopes
• 2.1.1(b) Atomic structure
• 2.1.1(c) Relative isotopic mass and relative atomic mass
• 2.1.1(d) Mass spectrometry
• 2.1.1(e) Relative molecular mass and relative formula mass

SPEC: 2.1.1(a) — Isotopes

You must know and apply:

• the meaning of isotopes;
• that isotopes are atoms of the same element with:
  • the same number of protons;
  • different numbers of neutrons;
  • different masses.

SPEC: 2.1.1(b) — Atomic structure

You must be able to determine:

• numbers of protons;
• numbers of neutrons;
• numbers of electrons;

for atoms and ions, given:

• atomic number;
• mass number;
• ionic charge.

You must know the relative:

• masses;
• charges;

of protons, neutrons and electrons.

SPEC: 2.1.1(c) — Relative mass

You must understand and explain:

• relative isotopic mass;
• relative atomic mass, (A_r);
• the use of carbon-12 as the atomic-mass standard;
• that relative atomic mass is a weighted mean.

SPEC: 2.1.1(d) — Mass spectrometry

You must know and apply the use of mass spectrometry in:

• determining relative isotopic masses;
• determining relative abundances of isotopes;
• calculating the relative atomic mass of an element from isotope data.

OCR limits calculations here to ions carrying a single charge.

Detailed construction/operation of a particular mass spectrometer is not required, but you should understand what a mass spectrum represents and how to use the data.

SPEC: 2.1.1(e) — Relative molecular mass and relative formula mass

You must use:

• relative molecular mass, (M_r) for simple molecules;
• relative formula mass for giant structures / ionic compounds;

and calculate these from relative atomic masses.

OCR does not require formal definitions of relative molecular mass or relative formula mass here, but you must use the terms correctly.

2. Core Content — Extreme Detail

2.1 SPEC: 2.1.1(a) — Isotopes

2.1.1 What is an isotope?

DEFINITION — Isotopes
Atoms of the same element that have the same number of protons but different numbers of neutrons, and therefore different masses.

This definition must be exact.

The two essential points are:

1. same number of protons;
2. different numbers of neutrons.

Do not define isotopes only as:

“Atoms with different mass numbers.”

That is incomplete because it does not establish that the atoms belong to the same element.

2.1.2 Why the proton number must be the same

The identity of an element is determined by its number of protons.

For example:

• 6 protons → carbon;
• 8 protons → oxygen;
• 11 protons → sodium;
• 17 protons → chlorine.

If the number of protons changed, the element would change.

Therefore isotopes must have the same proton number.

2.1.3 Why isotopes have different masses

The neutron has relative mass approximately:

[ 1 ]

so changing the number of neutrons changes the mass of the nucleus.

Example:

[ {}^{35}_{17}\mathrm{Cl} ]

and:

[ {}^{37}_{17}\mathrm{Cl} ]

both contain:

[ 17\text{ protons} ]

but:

[ {}^{35}\mathrm{Cl} ]

contains:

[ 18\text{ neutrons} ]

whereas:

[ {}^{37}\mathrm{Cl} ]

contains:

[ 20\text{ neutrons} ]

Therefore the isotopes have different masses.

2.1.4 Chemical properties of isotopes

Neutral isotopes of the same element have the same number of electrons.

Therefore they have the same electron arrangement.

Chemical reactions depend mainly on electrons, especially outer-shell electrons.

Hence isotopes of the same element have very similar chemical properties.

Examiner-safe phrasing

“The isotopes have the same number and arrangement of electrons, so they have the same chemical properties.”

Avoid:

“They react the same because they are the same element.”

That is a conclusion, not an explanation.

2.1.5 Physical properties of isotopes

Physical properties that depend on mass may differ between isotopes.

Examples include:

• rate of diffusion;
• density in some contexts;
• behaviour in a mass spectrometer.

This is because isotopes have different numbers of neutrons and therefore different masses.

2.2 SPEC: 2.1.1(b) — Atomic Structure

An atom contains:

• a central nucleus;
• electrons outside the nucleus.

The nucleus contains:

• protons;
• neutrons.

2.2.1 Proton

DEFINITION — Proton
A positively charged subatomic particle in the nucleus with relative charge (+1) and relative mass (1).

2.2.2 Neutron

DEFINITION — Neutron
An uncharged subatomic particle in the nucleus with relative charge (0) and relative mass (1).

2.2.3 Electron

DEFINITION — Electron
A negatively charged subatomic particle outside the nucleus with relative charge (-1) and very small relative mass, approximately (1/1836).

OCR commonly accepts the electron mass as negligible in many contexts, but the more precise A-level statement is:

[ \text{relative mass}\approx\frac{1}{1836} ]

2.2.4 Required particle data

2.3 Atomic Number and Mass Number

2.3.1 Atomic number

DEFINITION — Atomic number, (Z)
The number of protons in the nucleus of an atom.

The atomic number identifies the element.

For a neutral atom:

[ \boxed{ \text{electrons}=\text{protons}=Z } ]

2.3.2 Mass number

DEFINITION — Mass number, (A)
The total number of protons and neutrons in the nucleus.

Therefore:

[ \boxed{ A=p+n } ]

and:

[ \boxed{ n=A-Z } ]

2.3.3 Nuclide notation

A nuclide may be represented as:

[ {}^{A}_{Z}X ]

where:

• (X) = element symbol;
• (A) = mass number;
• (Z) = atomic number.

Example:

[ {}^{23}_{11}\mathrm{Na} ]

contains:

• 11 protons;
• (23-11=12) neutrons;
• 11 electrons if neutral.

2.4 Atoms and Ions

Ion formation changes the number of electrons.

It does not change the number of protons during ordinary chemical reactions.

2.4.1 Positive ions

A positive ion has lost electrons.

Example:

[ \mathrm{Mg}^{2+} ]

Magnesium has:

[ Z=12 ]

so:

• protons = 12;
• electrons = (12-2=10).

2.4.2 Negative ions

A negative ion has gained electrons.

Example:

[ \mathrm{Cl}^{-} ]

Chlorine has:

[ Z=17 ]

so:

• protons = 17;
• electrons = (17+1=18).

2.4.3 Universal charge relationship

For ions:

[ \boxed{ \text{ionic charge}

\text{protons}-\text{electrons} } ]

when the charge is expressed in units of the elementary charge.

2.4.4 Worked example

For:

[ {}^{52}_{24}\mathrm{Cr}^{3+} ]

Protons:

[ 24 ]

Neutrons:

[ 52-24=28 ]

Electrons:

[ 24-3=21 ]

Therefore:

[ \boxed{ 24p,\ 28n,\ 21e^- } ]

2.5 SPEC: 2.1.1(c) — Relative Isotopic Mass

Atoms are far too small for their masses to be conveniently expressed in grams.

Chemists therefore use a relative scale based on carbon-12.

2.5.1 Carbon-12 standard

The atomic-mass scale is defined relative to:

[ \frac{1}{12} ]

of the mass of one atom of:

[ {}^{12}\mathrm{C} ]

Carbon-12 is assigned exactly:

[ 12 ]

on this relative scale.

2.5.2 Relative isotopic mass

DEFINITION — Relative isotopic mass
The mass of an atom of an isotope compared with one-twelfth of the mass of an atom of carbon-12.

Relative isotopic mass:

• has no unit;
• may be close to, but is not necessarily exactly equal to, the mass number.

Important distinction

Mass number is an exact integer count of nucleons.

Relative isotopic mass is an experimentally measured relative mass.

2.6 Relative Atomic Mass

DEFINITION — Relative atomic mass, (A_r)
The weighted mean mass of an atom of an element compared with one-twelfth of the mass of an atom of carbon-12.

The word:

[ \boxed{\text{weighted}} ]

is essential.

This is because naturally occurring isotopes may not occur in equal abundances.

2.6.1 Why (A_r) is often not an integer

Chlorine has two abundant isotopes, approximately:

[ {}^{35}\mathrm{Cl} ]

and:

[ {}^{37}\mathrm{Cl} ]

with the lighter isotope more abundant.

Therefore the weighted mean is approximately:

[ 35.5 ]

rather than exactly 35 or 37.

2.6.2 Relative mass quantities have no unit

Relative isotopic mass and relative atomic mass are ratios.

Therefore:

[ \boxed{ A_r\text{ has no unit} } ]

Do not write:

[ 35.5\ \mathrm{g} ]

for the relative atomic mass of chlorine.

2.7 Calculating Relative Atomic Mass

General formula:

[ \boxed{ A_r= \frac{\sum(\text{isotopic mass}\times\text{relative abundance})} {\sum(\text{relative abundance})} } ]

If abundances are percentages totalling 100:

[ \boxed{ A_r= \frac{\sum(\text{isotopic mass}\times%\text{ abundance})}{100} } ]

2.7.1 Worked example using percentages

An element has:

• isotope 24: 79%;
• isotope 25: 10%;
• isotope 26: 11%.

[ A_r= \frac{(24\times79)+(25\times10)+(26\times11)}{100} ]

[

\frac{1896+250+286}{100} ]

[

24.32 ]

Therefore:

[ \boxed{ A_r=24.32 } ]

2.7.2 Worked example using arbitrary peak intensities

Mass spectrum:

• (m/z=10), intensity = 25;
• (m/z=11), intensity = 100.

Total abundance:

[ 25+100=125 ]

Weighted mean:

[ A_r= \frac{(10\times25)+(11\times100)}{125} ]

[

\frac{1350}{125} ]

[ \boxed{ A_r=10.8 } ]

Do not divide by 100 unless the values are percentages.

2.8 SPEC: 2.1.1(d) — Mass Spectrometry

OCR requires the use of mass spectrometry to determine:

• relative isotopic masses;
• relative abundances;
• relative atomic mass.

Detailed knowledge of the internal design of the instrument is not required in this specification point.

However, you should understand what the mass spectrum shows and why peaks correspond to isotopes.

2.8.1 What a mass spectrometer measures

A mass spectrometer separates/detects ions according to their:

[ \boxed{ m/z } ]

where:

• (m) = ion mass;
• (z) = ion charge.

OCR limits this section to ions with:

[ \boxed{ z=+1 } ]

Therefore, numerically:

[ m/z ]

corresponds directly to the relative isotopic mass of a monatomic ion.

2.8.2 Why ions are required

Mass spectrometry requires charged particles because ions can be controlled and detected using electric and/or magnetic effects.

Neutral atoms cannot be manipulated in the same way.

For a simple monatomic species, ionisation can be represented generically as:

[ X(g)\rightarrow X^+(g)+e^- ]

The exact instrument mechanism is not the focus of OCR 2.1.1(d).

2.8.3 The mass spectrum

DEFINITION — Mass spectrum
A graph showing ion signal/relative abundance against mass-to-charge ratio, (m/z).

Typical axes:

• x-axis: [ m/z ]
• y-axis: relative abundance or relative intensity.

Each isotope of an element can produce a separate peak.

2.8.4 Relative isotopic mass from a spectrum

For singly charged ions:

[ z=1 ]

so a peak at:

[ m/z=35 ]

corresponds approximately to isotope mass 35.

A second peak at:

[ m/z=37 ]

corresponds approximately to isotope mass 37.

2.8.5 Relative abundance from peak intensity

The relative height or area of each isotope peak indicates its relative abundance.

Example:

If the peaks have relative intensities:

[ 3:1 ]

then approximate abundances are:

[ 75%:25% ]

because:

[ \frac{3}{3+1}=0.75 ]

and:

[ \frac{1}{4}=0.25 ]

2.8.6 Calculating (A_r) from a mass spectrum

Suppose a spectrum contains:

• mass 35, relative abundance 75;
• mass 37, relative abundance 25.

Then:

[ A_r= \frac{(35\times75)+(37\times25)}{100} ]

[

\frac{2625+925}{100} ]

[

35.5 ]

Therefore:

[ \boxed{ A_r=35.5 } ]

2.8.7 Using peak ratios

Suppose two isotope peaks occur at:

[ 63,\ 65 ]

with intensity ratio:

[ 7:3 ]

Then:

[ A_r= \frac{(63\times7)+(65\times3)}{10} ]

[

\frac{441+195}{10} ]

[

63.6 ]

2.9 Identifying an Element from Isotope Data

A mass spectrum can be used to identify an element using:

• peak positions;
• relative abundance pattern;
• calculated (A_r).

Example:

A spectrum with peaks at:

[ 35,\ 37 ]

in approximately:

[ 3:1 ]

ratio is characteristic of chlorine.

Calculated:

[ A_r\approx35.5 ]

which supports identification.

2.10 Common Mass-Spectrometry Interpretation Rules

For OCR 2.1.1(d):

1. assume ions have a single positive charge unless told otherwise;
2. read isotope masses from (m/z);
3. use peak intensities as relative abundances;
4. calculate weighted mean for (A_r);
5. do not add unnecessary detail about instrument construction.

2.11 SPEC: 2.1.1(e) — Relative Molecular Mass

OCR uses relative molecular mass, (M_r) for substances made from discrete molecules.

Examples:

• (\mathrm{H_2O})
• (\mathrm{CO_2})
• (\mathrm{NH_3})
• (\mathrm{CH_4})

Although OCR does not require a formal definition here, the calculation is essential.

2.11.1 Calculating (M_r)

[ \boxed{ M_r=\sum A_r } ]

for all atoms in the molecular formula.

Example:

[ \mathrm{H_2O} ]

Using:

[ A_r(\mathrm H)=1.0 ]

and:

[ A_r(\mathrm O)=16.0 ]

[ M_r=(2\times1.0)+16.0 ]

[ \boxed{ M_r=18.0 } ]

No unit.

2.11.2 Example: carbon dioxide

[ \mathrm{CO_2} ]

[ M_r=12.0+(2\times16.0) ]

[ \boxed{ M_r=44.0 } ]

2.11.3 Example: sulfuric acid

[ \mathrm{H_2SO_4} ]

Using approximate values:

[ A_r(\mathrm H)=1.0 ]

[ A_r(\mathrm S)=32.1 ]

[ A_r(\mathrm O)=16.0 ]

[ M_r=(2\times1.0)+32.1+(4\times16.0) ]

[ \boxed{ M_r=98.1 } ]

2.12 Relative Formula Mass

For substances that do not exist as discrete molecules, OCR uses relative formula mass.

Examples:

• ionic lattices: [ \mathrm{NaCl} ]
• giant covalent structures: [ \mathrm{SiO_2} ]

The calculation is still the sum of the relative atomic masses in the formula unit.

2.12.1 Example: sodium chloride

[ \mathrm{NaCl} ]

[ A_r(\mathrm{Na})=23.0 ]

[ A_r(\mathrm{Cl})=35.5 ]

Therefore:

[ \text{relative formula mass}

23.0+35.5 ]

[ \boxed{ 58.5 } ]

2.12.2 Example: calcium nitrate

[ \mathrm{Ca(NO_3)_2} ]

[ =1(\mathrm{Ca})+2(\mathrm N)+6(\mathrm O) ]

Using:

[ 40.1+2(14.0)+6(16.0) ]

[ =40.1+28.0+96.0 ]

[ \boxed{ 164.1 } ]

The brackets matter because all atoms inside the bracket are multiplied by the external subscript.

2.13 Relative Mass vs Molar Mass

This distinction is essential.

Relative molecular/formula mass

Example:

[ M_r(\mathrm{H_2O})=18.0 ]

This has:

[ \boxed{\text{no unit}} ]

Molar mass

One mole of water has molar mass:

[ 18.0,\mathrm{g,mol^{-1}} ]

Molar mass is a dimensional quantity.

Do not write:

[ M_r=18.0,\mathrm{g,mol^{-1}} ]

That confuses relative mass with molar mass.

3. Exact OCR Exam Language

These are original examiner-safe structures modelled on OCR Chemistry A expectations.

3.1 Define

Isotope

“Atoms of the same element with the same number of protons but different numbers of neutrons.”

Relative isotopic mass

“The mass of an atom of an isotope compared with one-twelfth of the mass of an atom of carbon-12.”

Relative atomic mass

“The weighted mean mass of an atom of an element compared with one-twelfth of the mass of an atom of carbon-12.”

3.2 State

Give the fact directly.

Example:

“A proton has relative charge (+1) and relative mass 1.”

3.3 Calculate

Use:

1. correct formula;
2. substitution;
3. arithmetic;
4. appropriate precision;
5. no unit for relative masses.

3.4 Explain

Use:

structural difference → consequence

Example:

“The isotopes have different numbers of neutrons, so their masses differ.”

For chemical properties:

“They have the same number and arrangement of electrons, so they have the same chemical properties.”

3.5 Deduce

Use data to reach a justified conclusion.

Example:

“The peaks at (m/z=35) and (37) show that the element has two isotopes with those relative isotopic masses.”

3.6 Compare

Explicitly refer to both species.

“Both isotopes contain 17 protons, but chlorine-37 contains two more neutrons than chlorine-35.”

4. Common Misconceptions & Lost Marks

4.1 Isotopes = different numbers of protons

❌ Incorrect.

If proton number changes, the element changes.

✅ Isotopes have the same proton number and different neutron numbers.

4.2 “Same atomic number, different mass number” as the entire isotope definition

This may describe isotopes, but OCR's preferred chemistry wording is:

same number of protons, different numbers of neutrons.

Use the particle-level definition.

4.3 Mass number = relative atomic mass

❌ Incorrect.

Mass number:

[ A=p+n ]

for one nucleus.

Relative atomic mass:

[ A_r ]

is a weighted mean across naturally occurring isotopes.

4.4 Protons change when an ion forms

❌ Incorrect.

Chemical ion formation changes electrons.

4.5 Electron mass = exactly zero

The electron has a very small mass:

[ \approx\frac{1}{1836} ]

of a proton.

“Negligible” may be acceptable in some contexts, but “massless” is chemically false.

4.6 (A_r) has units

❌ Incorrect.

Relative masses are ratios.

No unit.

4.7 Using “weight” instead of mass

OCR mark schemes are strict about chemistry terminology.

Write:

“mass”

not:

“weight”

for relative mass definitions.

4.8 “Mass of an average atom”

Avoid this vague definition.

Use:

“weighted mean mass of an atom of an element…”

4.9 Forgetting the carbon-12 reference

For relative isotopic mass and (A_r), include:

[ \frac{1}{12} ]

of the mass of one carbon-12 atom.

4.10 Peak height = isotope mass

❌ Incorrect.

• x-position → (m/z);
• peak intensity → relative abundance.

4.11 Dividing by 100 when abundance data are not percentages

If values are relative intensities, divide by the total intensity.

4.12 Forgetting OCR's single-charge limitation

At this specification point, calculations are limited to singly charged ions.

So for monatomic ions:

[ m/z ]

maps directly to isotope mass.

4.13 Over-learning mass-spectrometer hardware

OCR explicitly focuses on the use of mass spectrometry here.

Do not waste revision time learning detailed engineering that is not required by this specification point.

4.14 (M_r) with units

❌:

[ M_r=44,\mathrm{g,mol^{-1}} ]

✅:

[ M_r=44 ]

Molar mass:

[ 44,\mathrm{g,mol^{-1}} ]

4.15 Using “relative molecular mass” for NaCl

NaCl is an ionic lattice, not a discrete molecule.

Use:

relative formula mass

4.16 Ignoring brackets in formula-mass calculations

For:

[ \mathrm{Ca(NO_3)_2} ]

the 2 multiplies both N and all three O atoms.

5. Worked Exam-Style Questions

Question 1 — Atomic Structure [4 marks]

For:

[ {}^{58}_{28}\mathrm{Ni}^{2+} ]

determine the number of:

(a) protons [1]

[ \boxed{28} ]

(b) neutrons [1]

[ 58-28=30 ]

[ \boxed{30} ]

(c) electrons [2]

Neutral Ni has 28 electrons.

A (2+) ion has lost two:

[ 28-2=26 ]

[ \boxed{26} ]

Question 2 — Isotopes [4 marks]

Explain why:

[ {}^{24}\mathrm{Mg} ]

and:

[ {}^{26}\mathrm{Mg} ]

are isotopes and why they have the same chemical properties.

Model answer

“They are isotopes because they are atoms of the same element with the same number of protons but different numbers of neutrons. Neutral atoms of the two isotopes have the same number and arrangement of electrons, so they have the same chemical properties.”

Indicative marks:

• same element / same protons;
• different neutrons;
• same electron number/configuration;
• same chemical properties linked to electron arrangement.

Question 3 — Relative Atomic Mass [5 marks]

An element has three isotopes:

Calculate (A_r).

Model answer

[ A_r= \frac{(24\times79)+(25\times10)+(26\times11)} {79+10+11} ]

[

\frac{2432}{100} ]

[ \boxed{ A_r=24.32 } ]

No unit.

Indicative marks:

• weighted-mean approach;
• correct isotope products;
• correct total abundance;
• correct numerical result;
• no unit / appropriate presentation.

Question 4 — Mass Spectrum [6 marks]

A mass spectrum of an element has two peaks:

• (m/z=63), relative intensity 69;
• (m/z=65), relative intensity 31.

(a) State what the two peak positions show. [2]

“The element contains isotopes with relative isotopic masses approximately 63 and 65.”

(b) Calculate the relative atomic mass. [4]

[ A_r= \frac{(63\times69)+(65\times31)}{100} ]

[

\frac{4347+2015}{100} ]

[

63.62 ]

[ \boxed{ A_r=63.62 } ]

Question 5 — Relative Formula Mass [5 marks]

Calculate the relative formula mass of:

[ \mathrm{Al_2(SO_4)_3} ]

using:

[ A_r(\mathrm{Al})=27.0 ]

[ A_r(\mathrm{S})=32.1 ]

[ A_r(\mathrm{O})=16.0 ]

Model answer

Al:

[ 2(27.0)=54.0 ]

S:

[ 3(32.1)=96.3 ]

O:

[ 12(16.0)=192.0 ]

Total:

[ 54.0+96.3+192.0 ]

[ \boxed{ 342.3 } ]

No unit.

Question 6 — A/A* Integrated Data Question [6 marks]

A student says:

“The atomic mass of chlorine is 35.5, so every chlorine atom has a mass number of 35.5.”

Evaluate this statement.

Full-mark model answer

“The statement is incorrect. Mass number is the total number of protons and neutrons in one nucleus and must therefore be a whole number. Chlorine occurs naturally as a mixture of isotopes, mainly chlorine-35 and chlorine-37. The value 35.5 is the relative atomic mass, which is a weighted mean of the relative isotopic masses according to their abundances. No individual chlorine nucleus has 35.5 nucleons.”

Indicative marking points:

• statement incorrect;
• mass number = protons + neutrons;
• mass number must be integer;
• chlorine has isotopes;
• (A_r) is weighted mean;
• no individual atom has mass number 35.5.

6. Links to Other Topics

6.1 2.1.2 — Compounds, Formulae and Equations

Particle counts and ionic charge support:

• writing ionic formulae;
• balancing charge;
• constructing chemical equations.

6.2 2.1.3 — Amount of Substance

Relative atomic mass and relative formula/molecular mass are essential for:

[ n=\frac{m}{M} ]

Molar mass numerically reflects the appropriate relative mass but has units:

[ \mathrm{g,mol^{-1}} ]

6.3 2.2.1 — Electronic Structure

OCR places:

• shells;
• subshells;
• orbitals;
• electron configurations;

in 2.2.1, not in this section.

This file must therefore not duplicate electronic-structure notes.

6.4 3.1 — Periodicity

Atomic identity and isotope concepts support interpretation of:

• periodic table position;
• nuclear charge;
• atomic trends.

6.5 4.2.4 — Analytical Techniques

Mass spectrometry later develops into:

• molecular-ion peaks;
• molecular mass;
• identification of organic compounds;
• M+1 peaks from carbon-13.

The isotope ideas introduced here become important again.

6.6 6.3 — Analysis

Advanced spectroscopy questions often combine:

• mass spectrometry;
• infrared spectroscopy;
• NMR;
• elemental composition.

7. Summary Notes

7.1 Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons and therefore different masses.

Same chemical properties because:

[ \boxed{ \text{same electron arrangement} } ]

7.2 Subatomic particles

7.3 Atomic number and mass number

[ \boxed{ Z=\text{protons} } ]

[ \boxed{ A=\text{protons}+\text{neutrons} } ]

[ \boxed{ \text{neutrons}=A-Z } ]

Neutral atom:

[ \boxed{ e^-=p } ]

Positive ion:

[ \boxed{ \text{electrons lost} } ]

Negative ion:

[ \boxed{ \text{electrons gained} } ]

7.4 Relative isotopic mass

Mass of one atom of an isotope compared with (1/12) of the mass of one carbon-12 atom.

No unit.

7.5 Relative atomic mass

Weighted mean mass of an atom of an element compared with (1/12) of the mass of one carbon-12 atom.

[ \boxed{ A_r= \frac{\sum(\text{isotopic mass}\times\text{abundance})} {\sum(\text{abundance})} } ]

No unit.

7.6 Mass spectrometry

OCR use:

• determine relative isotopic masses;
• determine relative abundances;
• calculate (A_r).

At this point:

[ \boxed{ \text{single-charge ions only} } ]

Mass spectrum:

• x-axis → (m/z);
• y-axis → relative abundance/intensity.

For (z=1):

[ \boxed{ m/z\approx\text{relative isotopic mass} } ]

Peak intensity:

[ \boxed{ \text{relative abundance} } ]

7.7 Relative molecular mass

For a simple molecule:

[ \boxed{ M_r=\sum A_r } ]

Example:

[ M_r(\mathrm{CO_2})=44.0 ]

No unit.

7.8 Relative formula mass

Use for:

• ionic lattices;
• giant structures.

Example:

[ \mathrm{NaCl} ]

[ 23.0+35.5=58.5 ]

No unit.

7.9 Relative mass vs molar mass

Relative mass:

[ \boxed{ \text{no unit} } ]

Molar mass:

[ \boxed{ \mathrm{g,mol^{-1}} } ]

Do not confuse them.