Collins Cambridge AS & A Level Chemistry by Collins

chapter

Atomic structure

Greylag geese travel thousands of miles across Europe, North Africa, Asia and into China. Avian flu, or bird flu, is an infectious disease that spreads among birds. In rare cases, it can infect humans. How the disease spreads from country to country is not clear, but we do know that migratory ducks and other water fowl can spread the disease. Scientists are using stable isotopes to understand the migratory patterns of these birds. The percentage abundance of isotopes such as carbon-13, hydrogen-2, sulfur-34 and strontium-87 is being measured in migratory birds and in different geographical environments. When the percentages match, the scientists can link the birds with the environment.

Prior understanding

You may remember from your previous course that atoms consist of protons, neutrons and electrons. You may recall the arrangement of electrons in an atom and how this links to the Periodic Table. You may have used the atomic number and mass number to describe the particles in an atom. You may be able to explain the existence of isotopes.

learning objectives

In this chapter, you will find out more about atomic structure, which will underpin many of the topics you study later. You will read how electrons are organised into shells and sub-shells and find out how to write full electronic configurations. You will explore the energy involved when atoms lose electrons (the ionisation energy) and the factors that influence it.

1.1 Particles in the atom and atomic radius (syllabus 1.1.1–1.1.7) 1.2 isotopes (syllabus 1.2.1–1.2.4) 1.3 electrons: energy levels and atomic orbitals (syllabus 1.3.1–1.3.9) 1.4 ionisation energy (syllabus 1.4.1–1.4.8)

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1 Atomic structure

1.1 Particles in the atom and atomic radius Atomic structure

Today, we know that atoms are mostly empty space. The central nucleus is very small and consists of protons and neutrons. Electrons are found in the empty space around the nucleus. Electrons are located in shells. Figure 1.1 shows a model of an atom.

protons neutrons

nucleus

electrons

The electrons are kept apart by their negative charge.

Nuclear binding forces allow protons to be close together.

Figure 1.1 This is a model of an atom. Atoms of different elements have different numbers of protons, neutrons and electrons.

At the beginning of the 20th century, scientists considered atoms to be negatively charged electrons in a sphere of positively charged material. Protons and neutrons had not been discovered. The model was known as ‘the plum pudding model’ (Figure 1.2). Rutherford and other scientists working in the UK carried out an experiment to fire positively charged alpha particles at gold foil, about 2000 atoms thick. Based on the ‘plum pudding’ model, the alpha particles should pass straight through the foil. It was like firing bullets at tissue paper. To their surprise, not all alpha particles passed straight through. Some were deflected and a few were completely deflected backwards. Figure 1.3 shows their experiment. many electrons with negative charge negative and positive charges cancel out spherical cloud of positive charge Figure 1.2 The ‘plum pudding’ model of the atom.

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1.1 Particles in the atom and atomic radius screen emits a flash when an alpha particle strikes it slightly deflected alpha particles

undeflected particles: most take this route

sometimes a particle is deflected through nearly 180°

beam of alpha particles

alpha particle source (radium)

lead shield to confine radiation Figure 1.3 Rutherford’s experiment: the deflection of alpha particles.

Rutherford concluded that:

gold foil about 2000 atoms thick

• The mass of the atom must be concentrated in a minute central region called the nucleus. • The nucleus must be positively charged in order to deflect the positively charged alpha particles. His findings are shown in Figure 1.4. Protons and neutrons were discovered later.

gold atom

alpha particle passes straight through

positively charged alpha particle

tiny positive nucleus ‘cloud’ of electrons

alpha particle is reflected Figure 1.4 Deflection of alpha particles by gold foil.

Protons, neutrons and electrons Experimental evidence has established the masses and charges of protons, neutrons and electrons. These are shown in Table 1.1. These are not distributed evenly over the atom. Most of the mass of the atom is in the nucleus which is positively charged. The electrons are outside the nucleus and are negatively charged. They have smaller mass than protons and neutrons, and the mass of the electrons can usually be ignored.

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1 Atomic structure Particle

Mass/kg

Charge/C

Relative mass

Relative charge

Electron

9.109 × 10−31

1.602 × 10−19

5.45 × 10−4

−1

Proton

1.672 × 10

1.602 × 10

Neutron

1.674 × 10−27

−27

−19

Note: The mass of the electron is so small compared to the mass of the proton and neutron that chemists often take it to be zero. Table 1.1 The mass and charge of protons, neutrons and electrons.

Because the values for mass are so small, the idea of relative mass is used. The relative mass of a proton is 1 and that of a neutron is 1. The relative mass of the electron is 5.45 × 10 –4 or 1 . This means that an electron has a mass 5.45 × 10 –4 the mass of a 1837 proton. Charges on sub-atomic particles are also given relative to one another. A proton has a relative charge of +1 and an electron has a relative charge of −1. A neutron has no charge. The protons and neutrons together are called nucleons. Protons and neutrons in the nucleus are held together by the strong nuclear force that acts over the small nucleus and binds the nucleons together. As atoms of any element are neutral, the number of protons (positive charge) must equal the number of electrons (negative charge). The atoms of all elements, except hydrogen, contain these three fundamental particles.

Sub-atomic particles in an electric field

Experiments with protons, neutrons and electrons in an electric field confirm their relative charges. Consider a beam of protons, neutrons and electrons, all of which are moving at the same velocity when they enter the space between two electrically charged plates. One plate is positively charged, the other is negatively charged, creating an electric field in the space between them. Remember that opposite charges attract and like charges repel. Protons are positively charged and will be deflected away from the positive plate and towards the negative plate. Neutrons have no charge and will continue moving in a straight line. Electrons are negatively charged and will be deflected away from the negative plate and towards the positive plate. Electrons have a much lower mass than protons. If the electrons and protons are travelling at the same velocity, then the electrons will be deflected more strongly than the heavier protons. Figure 1.5 shows the deflections. positive plate beam of subatomic particles

electrons neutrons protons negative plate

Figure 1.5 Protons, neutrons and electrons in an electric field.

The experiment shows that: • • • •

Protons are positively charged. Electrons are negatively charged. Neutrons have no charge. Electrons have a very small mass compared to protons.

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1.1 Particles in the atom and atomic radius

Assignment 1.1: Size, atoms and sub-atomic particles

A single carbon atom measures about one ten-billionth of a metre across, a dimension so small that it is impossible to imagine. (This is the US billion; 1 billion = 1000 million.) The nucleus is a thousand times smaller again, and the electron a hundred thousand times smaller than that!

Figure 1.6 A single carbon atom measures about one ten-billionth of a metre across.

Because the numbers are so unimaginably small, scientists do not use grams and metres to describe atoms and sub-atomic particles. They use a different set of units. You have already come across the idea of relative masses. Protons and neutrons both have a relative mass of 1. We say these have a mass of 1 relative mass unit. The electron is a mere 0.000 545 relative mass units. Clearly, even with relative masses you have some awkward numbers.

Tip Practise standard form calculations on your calculator if you need to. You will need to be confident with these later in the course.

Tip Do not use more significant figures in your answer than are used in the question.

A1. Assuming a carbon atom has a diameter of one ten-billionth of a metre, what is its diameter in metres in standard form? A2. What is the diameter in metres of an electron in standard form? A3. a. An atom of hydrogen contains only a proton and an electron. Calculate the mass of the hydrogen atom in kilograms. b. A molecule of hydrogen contains two atoms. Calculate the mass of a hydrogen molecule in grams. c. How many electrons have the same mass as a single neutron? A4. Convert these quantities into measurements in grams, expressed in standard form: a. the mass of a neutron b. 200 million electrons c. 10 gold coins weighing a total of 0.311 kg. A5. A uranium atom contains 92 electrons. Calculate the mass, in kilograms, of protons in the atom. A6. How many times heavier is the nucleus of a helium atom (two protons and two neutrons) than its electrons?

Atomic number and mass number Different elements have different numbers of electrons, protons and neutrons in their atoms. It is the number of protons in the nucleus of an atom that identifies the element. An atom may gain or transfer electrons to become an ion, but it is still an ion of the same element. A hydrogen atom has one proton and one electron. If a hydrogen atom loses its electron it forms the H+ ion. If it gains an electron it forms the H– ion, but it is still hydrogen. Similarly, an atom of an element can also have one or two more or fewer neutrons and still remain the same element. These are isotopes, like the ones used to track migrating birds. An atom can be defined using two numbers: the atomic number and the mass number. 5

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1 Atomic structure Atomic number The atomic number is also called the proton number. It has the symbol Z. The atomic number of an element is the number of protons in the nucleus of the atom. Since atoms are neutral and a proton has a relative charge of +1, the number of protons in the nucleus equals the number of electrons outside the nucleus. All atoms of the same element have the same atomic number.

Mass number The mass number is also called the nucleon number. It has the symbol A. The mass number of an element is the total number of protons and neutrons in the nucleus of an atom. It is a measure of its mass compared with other types of atom. Even in heavy atoms, the electron’s mass is so small that it makes little difference to the overall mass of the atom. Protons and neutrons both have a mass of 1, so:

mass number (A) = number of protons (Z) + number of neutrons (n)

A=Z+n You can calculate the number of neutrons in the nucleus using:

number of neutrons (n) = mass number (A) − atomic number (Z). Worked example

How many neutrons are present in an atom with 79 protons and a mass number of 197? Answer Since number of neutrons (n) = mass number (A) − atomic number (Z), n = 197 − 79 = 118 neutrons

The following notation is used to show the atomic number and mass number of an atom: mass number

atomic number

23 11

Figure 1.7 Notation used to show the atomic number and mass number of an atom.

1. How many protons, neutrons and electrons do the following atoms and ions have? a. an element with mass number 19 and atomic number 9 b. an element with mass number 210 and atomic number 85 c. an ion with one positive charge, mass number 23 and atomic number 11 d. an ion with three negative charges, mass number 31 and atomic number 15 e. an ion with three positive charges, mass number 52 and atomic number 24. 2. Copy and complete the table for the following particles. The first one has been done for you. Number of protons

Number of neutrons

Number of electrons

Symbol

P3–

Table 1.2

3. A beam of protons is fired between two electrically charged plates. Describe what will happen and why. 6

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1.1 Particles in the atom and atomic radius

Atomic radius and ionic radius

Atomic radius Atoms are extremely small. Rutherford first determined the size of an atom. He found that gold atoms have a radius of about 1 × 10 –10 m. He defined atomic radius as the distance from the centre of the atom to its outermost electrons, as shown in Figure 1.8. (It is actually very difficult to decide where the edge of the electron cloud is, so this distance is very approximate.)

atomic radius

Figure 1.8 The atomic radius of an atom.

metallic radius (measured in nm)

covalent radius (measured in nm)

The term atomic radius is used to measure the size of an atom, but the size of an atom depends on the space occupied by the electrons. The electron cloud of an atom has no fixed dimensions, because an atom’s electrons could be found anywhere. When you use shells to locate an electron, one of these is the most likely place to find the electron. Chemists talk about the probability of the whereabouts of an electron. It may be elsewhere, which makes identifying the edge of an atom a problem. Atomic radius is usually considered to be half the shortest distance from the nucleus of one atom to that of the nearest atom. Because atoms combine with different types of bonding, there are different types of atomic radii. Covalent radius and metallic radius are two types of atomic radius (Figure 1.9).

bonding pair of electrons

The covalent radius is half the shortest distance from one nucleus to the next in a covalently bonded molecule.

metal ion

The metallic radius is half the shortest distance from one metal ion nucleus to the next.

Figure 1.9 Covalent and metallic radii. ‘nm’ is the abbreviation for nanometre. It is 1 × 10–9 m.

The nucleus in an atom is very small. The atom’s size depends mostly on the space its electrons occupy. You might expect that, as the number of electrons increases, so do the size of the atom and the atomic radius. This is true for atoms of elements in the same group. Atomic radius increases down the group because an extra shell of electrons is added with each successive element. You can see this trend in Figure 1.10. Metallic radii are given for the metals. Covalent radii are given for non-metals.

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1 Atomic structure Group 1

0.152

0.112

0.080

0.077

0.074

0.066

0.064

0.186

0.160

0.143

0.117

0.110

0.104

0.099

0.231

0.197

0.122

0.121

0.116

0.114

0.244

0.215

0.162

0.141

0.137

0.133

0.262

0.217

0.171

17 F

Period

0.175

0.146

0.14

0.140

Figure 1.10 The atomic radii (nm) of some elements in the Periodic Table.

You can revisit the section on ‘protons, neutrons and electrons’ in Section 1.1 to remind yourself of the strong nuclear charge that keeps the nucleons together.

But this trend is not repeated across a period. In fact, moving across a period the atomic radius decreases. In Period 3 all of the outer electrons are in the third shell or energy level, n = 3. There are two inner shells of electrons shielding the nuclear charge. Moving across the period, a proton is added to the nucleus of each element. An electron is also added, but it goes into the same shell and the positive charge of the nucleus acts over a definite area. The increasing positive nuclear charge pulls with greater force on the negative electrons. All the electron shells are drawn slightly closer to the nucleus. The overall effect is to pull all electrons closer to the nucleus. For example, a sodium atom is larger than a magnesium atom because the nuclear charge of magnesium is greater than that of sodium. Both have their outer electrons in the same shell, with the same amount of shielding by the full inner electron shells. But, the electrons in magnesium are pulled closer to the nucleus than they are in sodium, resulting in a slightly smaller radius. For the same reason, magnesium is larger than aluminium. You can see the same trend across all the periods in Figure 1.10.

Link

Ionic radius Ionic radii are difficult to measure. They are usually considered to be the radius of an ion in a crystal lattice. Figure 1.11 shows the ionic radii of some elements in the Periodic Table.

Li+

Be2+

B3+

C4+

N3–

O2–

Na+

Mg2+

Al3+

Si4+ 0.041

0.212

0.184

K+

Ca2+

Ga3+

Ge4+

As3–

Se2–

Rb+

Sr2+

In3+

Sn4+

Sb5+

Te2–

Cs+

Ba2+

Tl3+

Pb4+

Bi5+

0.060

0.095

0.138

0.152

0.167

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0.031

0.065

0.099

0.112

0.135

0.020

0.050

0.062

0.081

0.095

0.015

0.093

0.112

0.120

0.171

P3–

0.222

0.062

0.140

S2–

0.198

0.221

F–

0.136

Cl–

0.181

Br–

0.195

I–

0.216

0.074

Figure 1.11 The ionic radii (nm) of elements in the Periodic Table. The spheres show the relative sizes of the ions.

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1.2 Isotopes

Because ionic radii are difficult to measure, you may find different values in different books and websites. Only compare values from the same source.

4. Explain the following: a. Potassium has a larger atomic radius than calcium. b. Lithium has a smaller atomic radius than potassium. c. It is more convenient to measure atomic radii in nanometres than metres. d. Atomic radii increase down Group 1. 5. Explain why: a. A sodium atom has a larger radius than a sodium ion. b. A nitrogen ion has a larger radius than a lithium ion. c. A potassium ion has a larger radius than a sodium ion. d. A sodium ion has a larger radius than a magnesium ion. e. A chloride ion has a larger radius than a chlorine atom.

Tip

As you might expect, the ionic radius increases down a group as extra shells of electrons are added. The trends across a period depend on whether a positive ion or a negative ion is formed. When positive ions form, electrons are transferred and the ion has the electronic configuration of the preceding noble gas. An electron shell is lost. The number of protons in the nucleus continues to increase across the period, increasing the nuclear charge. This increasing nuclear charge pulls the electrons more strongly and the ionic radius decreases. When negative ions form, extra electrons are added to the outer shell and the ionic radius increases. The extra electrons are held less tightly than in the neutral atom since there is no change in the number of protons in the nucleus, and also the repulsive forces in the outer shell are increased.

Key ideas

➜➜Atoms have a small dense nucleus consisting of protons and neutrons with electrons in the empty space around the nucleus. ➜➜Protons and neutrons have a relative mass of 1; electrons have a relative mass of 5.45 × 10 –4 and their mass can usually be ignored. ➜➜Protons have a relative charge of +1, electrons have a relative charge of –1 and neutrons have no charge. ➜➜Atomic (proton) number (Z) is the number of protons in one atom. ➜➜Mass number (A) = number of protons + the number of neutrons. ➜➜The number of neutrons in an atom = mass number − atomic number. ➜➜Atomic radii increase down a group because the number of shells increases. ➜➜Atomic radii decrease across a period because the positive nuclear charge increases.

1.2 Isotopes

All atoms of the same element have the same number of protons and the same atomic number. However, atoms of the same element may have different numbers of neutrons and different mass numbers. These are called isotopes. Isotopes have the same atomic number, but different mass numbers. Isotopes can be identified using the notation yxA, where x = the mass number, y = the atomic number and A is the chemical symbol for the element. Most elements have isotopes. Carbon has three isotopes. Table 1.3 shows carbon’s three isotopes.

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1 Atomic structure Name

Notation

Number Number of Number of Atomic Mass Abundance/% of protons neutrons electrons number number

98.9

1.1

1 × 10 –10

carbon-12

12 6

carbon-13

13 6

carbon-14

14 6

Table 1.3 Carbon isotopes.

Properties of isotopes

Chemical properties The chemical properties of an element depend on the number and arrangement of the electrons in its atoms. It is the electrons that are rearranged when a chemical reaction takes place. Since all isotopes of an element have the same number and arrangement of electrons, they also have the same chemical properties.

You will find out how mass spectrometry is used to detect isotopes in Chapter 21.

Carbon-12 and carbon-13 are stable isotopes. Their nuclei do not break up spontaneously. Carbon-14 is not stable and breaks up, releasing radioactivity and forming a different element.

Link

Physical properties Many of the physical properties of an element depend on its mass. Since isotopes of the same element have different mass numbers, they will also have slightly different physical properties. As mass (g) density = volume (cm3 )

different isotopes of the same element will have different densities. Boiling point, melting point and rate of diffusion are also dependent on an element’s mass. 12 g of carbon-12 and 13 g of carbon-13 contain the same number of atoms and occupy the same volume. Carbon-13 will be denser than carbon-12.

Two uses of isotopes 1. Uranium is used as a fuel in the atomic energy industry. 238 235 Uranium has two major isotopes, 235 92U and 92U. Isotope 92U is 0.28% abundant 238 and 92U is 99.71% abundant. Samples of uranium contain both isotopes, but only uranium-235 can be used by the atomic energy industry. The two isotopes therefore have to be separated. The stages are: • Uranium is reacted with hydrogen fluoride (HF) to produce uranium hexafluoride: 235 238 92UF6 and 92UF6. It is easier to convert uranium hexafluoride to a gas than uranium. • Uranium hexafluoride is vaporised and allowed to diffuse. 235 235 • 238 92UF6 diffuses more slowly than 92UF6 because it is heavier. The isotopes 92UF6 238 and 92UF6 can be collected separately. 2. Carbon-14 is used to date remains of living things. Carbon-14 is an unstable isotope of carbon. It has a half-life of about 5600 years. Carbon-14 is continually made in the upper atmosphere, and the percentage abundance of carbon-14 stays constant on Earth. Living things take in carbon during their lives, including some carbon-14. When they die, they stop taking in carbon and the carbon-14 nuclei in their tissues decay. After 5600 years, half the original amount will remain. The ratio of carbon-12 to carbon-14 in the remains of living things is used to determine the age of the remains (Figure 1.12).

Figure 1.12 This skull can be dated by measuring the ratio of carbon-12 to carbon-14 it contains. The technique is called radiocarbon dating.

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6. Nitrogen has two stable isotopes, 147N and 157N. State the number of protons, neutrons and electrons in each isotope. 7. Hydrogen has three isotopes, 11H, 21H and 31H. State the number of protons, neutrons and electrons in each isotope.

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1.3 Electrons: energy levels and atomic orbitals 8. Explain how the separation of the two major isotopes of uranium relies on their different mass numbers. 37 9. Chlorine has two stable isotopes, 35 17Cl and 17Cl. Chlorine gas is injected into one end of a long tube and allowed to diffuse. Which isotope will reach the other end first and why?

1.3 Electrons: energy levels and atomic orbitals

➜➜Isotopes of an element have the same number of protons and electrons in their atoms, but different numbers of neutrons. ➜➜Isotopes of an element have the same atomic number but different mass numbers. ➜➜Isotopes of the same element have the same chemical properties. ➜➜Isotopes of the same element have different physical properties.

Key ideas

Shells and sub-shells

Max no. of sub-shells

Max no. electrons in sub-shell

Max. no. electrons in main shell

s p

1 3

2 6

s p d

1 3 5

2 6 10

s p d f

1 3 5 7

2 6 10 14

Sub-shell

Main shell

Electrons are arranged in electron shells around the nucleus. Electrons can be described as being in a particular shell. The energy associated with electrons in each shell increases for successive shells. Electron shells are also referred to as energy levels. Each energy level is given a number called the principal quantum number (n). The first shell is the n = 1 energy level and is closest to the nucleus. The second shell is the n = 2 energy level and is the next shell, and so on. Each additional electron occupies the lowest possible energy level. This is called the ground state of the electron. A hydrogen atom has one electron, which occupies the n = 1 energy level. The electron is in its ground state. Within each shell, there are sub-shells, or orbitals. The number of sub-shells in each shell is shown in Table 1.4.

Table 1.4 Shells, sub-shells and numbers of electrons.

The sub-shells are given the letters s, p, d and f. Table 1.4 shows that the first shell has a maximum of two electrons and that they are both in sub-shell s. The second electron shell has a maximum of eight electrons, two of which are in sub-shell s and six in sub-shell p. This sequence of sub-shells also corresponds to an increase in energy. The electrons in each sub-shell have a definite amount of energy. Electrons cannot have an amount 11

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1 Atomic structure of energy between the energies of the sub-shells. This is shown in Figure 1.13. Each additional electron goes into the sub-shell with the next lowest energy. The order of filling is the same as the order of the elements in the Periodic Table. 3d e n e r g y

1s Figure 1.13 The energies of the sub-shells in an atom with many electrons.

Figure 1.13 shows the increasing energy of the sub-shells. You may notice that t he 4s sub-shell has lower energy than the 3d sub-shell. This means that electrons occupy the 4s sub-shell before the 3d sub-shell. You will read more about this in the section on electronic configurations.

Orbitals Electrons are constantly moving, and it is impossible to know the exact position of an electron at any given time. However, measurements of the density of electrons as they move round the nucleus show that there are regions where it is highly probable to find an electron. These regions of high probability are called orbitals. Each s, p, d and f sub-shell corresponds to a differently shaped orbital. The s orbitals are spherical. The p orbitals have two lobes. The shapes of s and p orbitals are shown in Figure 1.14. Each orbital can hold two electrons, which spin in opposite directions. Table 1.5 shows the numbers of electrons and orbitals in the sub-shells.

s orbitals

3s z

p orbitals

p x orbital

the three p orbitals combined z pz

p y orbital

px x py y

p z orbital

Figure 1.14 The three-dimensional shapes of the s and p orbitals.

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1.3 Electrons: energy levels and atomic orbitals Sub-level

Number of orbitals in sub-shell

Maximum number of electrons

Table 1.5 The maximum number of electrons per orbital.

The arrangement of electrons in an atom can be written as symbols in an electronic configuration. The electronic configuration includes sub-shells as well as shells, and shows the number of electrons in each.

• Hydrogen has one electron in shell 1, sub-shell s. Its electronic configuration is 1s. • Helium has two electrons with opposite spin. Its electronic configuration is 1s2. • Lithium has two electrons in 1s and one in 2s. Its electronic configuration is 1s22s1.

Electronic configuration of atoms

1s 1s

3s2

Electronic configuration

Shells/sub-shells 2p

2s 2

You can also draw a diagram showing the electrons in boxes, or a spin diagram for each sub-shell. These show the direction of spin of all the electrons (Table 1.6). So, you can represent the 12 electrons in the shells/ sub-shells of magnesium in two ways: electronic configuration or a spin diagram.

Spin diagram

Table 1.6 Spin diagram for subshells of magnesium.

Link You will read more about the electronic configuration of the transition elements in Chapter 27.

The electronic configuration for magnesium is: 1s22s22p63s2. Between hydrogen and argon, electrons of increasing energy are added, one per element, in sub-shell order 1s, 2s, 2p, 3s, 3p. Then, for potassium, the next electron skips sub-shell 3d and goes into 4s. Although shell 3 energies are lower overall than shell 4 energies, the 3d sub-shell has a higher energy than the 4s sub-shell, as shown in Figure 1.13. The order of filling is the same as the order of elements in the Periodic Table, and sub-shell 4s is filled before 3d. Later, you will see that the chemical properties of elements reflect the energy levels of electrons.

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1 Atomic structure Z Element

Electron configuration

Electron spin diagram

1s H

1s 1

1s 2

1s 2 2s 1

1s 2 2s 2

1s 2 2s 2 2p 1

1s 2 2s 2 2p 2

1s 2 2s 2 2p 3

1s 2 2s 2 2p 4

1s 2 2s 2 2p 5

1s 2 2s 2 2p 6

1s 2 2s 2 2p 6 3s 1

1s 2 2s 2 2p 6 3s 2

1s 2 2s 2 2p 6 3s 2 3p 1

1s 2 2s 2 2p 6 3s 2 3p 2

1s 2 2s 2 2p 6 3s 2 3p 3

1s 2 2s 2 2p 6 3s 2 3p 4

1s 2 2s 2 2p 6 3s 2 3p 5

1s 2 2s 2 2p 6 3s 2 3p 6

1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1

1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 8 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 2

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6

Table 1.7 Electronic configurations and spin diagrams for the first 36 elements.

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1.3 Electrons: energy levels and atomic orbitals

Filling orbitals You have seen that the arrows in electron spin diagrams indicate the direction of spin of the electrons and whether there are one or two electrons per orbital. The electrons fill the orbitals in a set order. As electrons are all negatively charged, they repel each other. Electrons organise themselves so that they remain unpaired and fill the maximum number of sub-shells possible. As you have seen, for the p sub-shells this means that electrons first occupy empty orbitals that have the same energy and are parallel spinned. When these orbitals each have one electron, additional electrons are spin-paired; the second electron in an orbital will spin in the opposite direction. This is how the 2p orbitals fill: Electron 2p1 in boron is:

Electrons 2p2 in carbon are:

Electrons 2p3 in nitrogen are:

Electrons 2p5 in fluorine are:

There is now one electron in each orbital. The next electron goes into the first orbital and spins in the opposite direction, so that: Electrons 2p4 in oxygen are:

Electrons 2p6 in neon are:

Atoms with one or more unpaired electrons are called free radicals. In Table 1.7, shell 1 in helium is filled. The next element with a filled level is neon, which has the electronic configuration 1s22s22p6. Since the outermost shell is complete, these elements are very stable. They are known as the noble gases. Noble gas configurations are used to write shorthand electronic configurations. For example, the full electronic configuration for potassium is 1s22s22p63s23p64s1. The shorthand form is [Ar]4s1. Similarly, the shorthand electronic configuration for phosphorus is [Ne]3s23p3. Having eight electrons in the outer shells of atoms in compounds is very stable. This is the octet rule.

Tip When writing shorthand electronic configurations, use the noble gas that immediately precedes the element.

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1 Atomic structure Worked example Write the full electronic configuration for silicon. Answer Firstly, use your Periodic Table to find the number of electrons in one atom of silicon. An atom of silicon has 14 electrons. The n = 1 energy level has two electrons: 1s2. The n = 2 energy level has eight electrons: 2s22p6. The n = 3 energy level has four electrons are 3s23p2. The electronic configuration of silicon is 1s22s22p63s23p2.

When writing electronic configurations, check that the number of electrons (the superscripts) add up to the number of electrons you are representing.

Tip

Electronic configuration of ions

An ion is an atom in which either:

• one or more electrons have been removed, producing a positively charged ion, or • one or more electrons have been added, producing a negatively charged ion.

Worked example

Note that the maximum number of electrons that can be removed or added to form ions for Periods 1 to 3 is three.

What is the electronic configuration of the sodium ion, Na+?

Answer The electronic configuration of the sodium atom is: 1s22s22p63s1. In Na+ the outermost electron, 3s1, has been removed. This is the electron of highest energy in sodium, and so takes the least energy to remove.

The electronic configuration of Na+ is 1s22s22p6.

10. What do 2, p and the superscript 6 represent in 2p6? 11. Use Table 1.5 to help you write shorthand electronic configurations for: a. sulfur e. silicon b. aluminium f. iron c. calcium g. bromine d. scandium h. zinc. 12. Explain why a boron atom is a free radical but a beryllium atom is not. 13. Write the full electronic configuration for: d. Br– a. Ca2+ – b. Cl e. N3– c. Al3+ 14. Write the electronic configurations for atoms and ions with the following atomic numbers and charges: a. atomic number 9, charge = –1 b. atomic number 34, no charge c. atomic number 16, charge = –2 d. atomic number 13, charge = +3 e. atomic number 36, no charge.

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1.4 Ionisation energy

➜➜Electrons are arranged in shells and sub-shells. ➜➜Electrons in sub-shells are located in orbitals (s, p, d and f). ➜➜The energy of successive shells and sub-shells increases with distance from the nucleus. ➜➜The principal quantum number, n, is used to number the energy levels. ➜➜n = 1 is the energy level closest to the nucleus. ➜➜Two electrons can occupy the s sub-shells, six electrons can occupy the p sub-shells and ten electrons can occupy the d sub-shells. ➜➜Electronic configurations show the number of electrons in each sub-shell. ➜➜‘Electron in boxes’ diagrams show which orbitals are occupied in a sub-shell and the direction of spin of the electrons.

Key ideas

The energy required to remove an electron completely from an atom in its gaseous state is called the ionisation energy, IE. The energy required to remove the first electron is called the first ionisation energy and can be written as: M(g) → M+(g) + e– The energy required to remove the second electron from an atom is called the second ionisation energy and can be written as: M+(g) → M2+(g) + e– Ionisation energy values for removing the second and subsequent electrons are called successive ionisation energies. The ionisation energy for one atom is so small that, for convenience, ionisation energies are measured per mole of atoms, in kJ mol–1. The first ionisation energy is the enthalpy change (energy change) when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge.

Trends in ionisation energy

First ionisation energy/kJ mol–1

Different elements have different ionisation energies. For example, 418 kJ are required to remove one mole of electrons from one mole of gaseous potassium atoms, but 1260 kJ are required to remove one mole of electrons from one mole of gaseous chlorine atoms. Ionisation energies depend on the strength of the attraction between the outer electron and the nucleus. The outer electron in a chlorine atom is attracted to the nucleus more strongly than the outer electron in a potassium atom. Ionisation energies have been calculated for all but a few of the very heavy elements in the Periodic Table. Figure 1.15 shows the first ionisation energies for the elements from hydrogen to caesium plotted against their atomic number. He

2500

2000

1000

P Mg

500 0

1500

Li 0

Na 5

Zn As Fe Ni Se Ca Ti Cr Ge Mn Co Cu V Ga Sc K 20

Cs 40

Atomic number Figure 1.15 The first ionisation energies for the elements from hydrogen to caesium.

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1 Atomic structure The patterns in Figure 1.15 can be explained by: • • • •

the nuclear charge the atomic or ionic radius the shielding of the nuclear charge by inner shells and sub-shells of electrons the spin-pair repulsion when two electrons spin in opposite directions in one orbital.

First ionisation energy trends down a group Electron from sub-level: 2s

800

700 4s

600

500 400 Be

Mg Ca Sr Group 2 elements

300

Firstionisation energy/kJ mol–1

900

Figure 1.16 The first ionisation energies of Group 2 elements.

The Group 2 elements, beryllium to barium, are reactive metals. The outer shells of these elements contain a pair of electrons in an s orbital. The first ionisation energy measures how much energy is needed to remove one mole of these electrons from one mole of atoms. Figure 1.16 shows how the first ionisation energies decrease down Group 2. That means that the first electron becomes easier to remove, even though the nuclear charge is increasing. This is because:

• the number of electron shells between the outer electron and the nucleus is increasing; the electron shells shield the outer electron from the nuclear charge, and • the atomic radius is increasing as you go down Group 2; the distance between the outer electron and the nucleus is increasing. So, the outer electrons are easier to remove and the first ionisation energies decrease.

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First ionisation energy/kJ mol–1

First ionisation energy trends across a period There are similar trends in first ionisation energy across Periods 2 and 3 in the Periodic Table. Figure 1.17 shows the trends for Period 3. 2200 Electron from sub-level: 2p 2000 1800 1600 1400 1200 3p 1000 3p 3p 800 3p 3s 600 3p 400 3s 200 0 Ne Na Mg Al Si P S Cl Period 3 elements (plus Ne and K)

4s Ar

Figure 1.17 First ionisation energies of Period 3 elements from sodium to argon.

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1.4 Ionisation energy

• The outermost electron for sodium enters the 3s orbital. The electrons in the first and second shells shield the 3s electron from the positive charge of the nucleus and it is relatively easy to remove. • Magnesium has one more electron than sodium, and this completes the 3s orbital and spins in the opposite direction. Magnesium also has an extra proton, so the nuclear charge has increased. More energy is needed to remove magnesium’s first electron. Magnesium’s first ionisation energy is higher than sodium’s. • Aluminium has one more electron than magnesium and this is the first to enter a 3p orbital. p orbitals have higher energy than s orbitals. Aluminium’s first electron is easier to remove than the 3s electron of magnesium. The first ionisation energy drops. • The extra electron that silicon has compared to aluminium, and that phosphorus has compared to silicon, fills the remaining empty 3p orbitals. At the same time, the nuclear charge is increasing and more energy is needed to remove these electrons. The first ionisation energies increase from aluminium to phosphorus. • Sulfur’s first electron enters a 3p orbital already containing one electron. These spin in opposite directions and repel each other. It takes less energy to remove the first electron from sulfur than to remove the first electron from phosphorus. Its first ionisation energy is lower. The electrons for chlorine and argon fill the remaining 3p orbitals. The nuclear charge continues to increase and the first ionisation energy increases as more energy is needed to remove an electron.

As you move across Period 3, each element has one more electron than the last. This electron fills the first available empty orbital. Overall, first ionisation energies increase across Period 3 because the nuclear charge is increasing and electrons are attracted more strongly. These are the trends across Period 3:

Successive ionisation energies An element has as many successive ionisation energies as it has electrons. Magnesium has the electronic configuration 2,8,2. Its first three successive ionisation energies are: Mg(g) → Mg+(g) + e– first ionisation energy = +738 kJ mol–1 Mg+(g) → Mg2+(g) + e– second ionisation energy = +1451 kJ mol–1

Mg2+(g) → Mg3+(g) + e– third ionisation energy = +7733 kJ mol–1 Figure 1.18 shows a graph of log10 of the successive ionisation energy against the number of electrons removed for magnesium. We use log10 of the ionisation energies to make the numbers easier to handle. (Remember that adding 1 to a log10 of a number is the same as multiplying the number by 10, so the log10 increases much more slowly than its number.)

Link You can find out more about logarithms in Chapter 24.

log10 (successive ionisation energy)

6 5

Tip

You may need to talk to your maths teacher if you have not studied logs before.

3 2 1 0

2 4 6 8 10 12 Number of electrons removed

Figure 1.18 The successive ionisation energies of magnesium.

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1 Atomic structure

2s electrons

unpaired 2p electrons

Ionisation energy (kJ mol–1)

38000 36000 34000 32000 30000 28000 26000 24000 22000 20000 18000 16000 14000 12000 10000 8000 6000 4000 2000 0

The first two electrons, the 3s2 electrons, are removed from the third shell. The second ionisation energy is higher than the first ionisation energy. This is because the nuclear charge stays the same − there are still 12 protons in the nucleus, and each time an electron is removed the remaining electrons are attracted more strongly. The large increase between the second and third electron is because the third electron is taken from an inner shell, the second shell. The ionic radius has decreased, the third electron is closer to the nucleus and there is one fewer shell to shield the nuclear charge. A large increase in energy is required to remove this 2p electron. This pattern is repeated for the eleventh ionisation energy because the electron is removed from the first shell. The increase from the third to the tenth electron shows electrons being removed from the second shell. The nuclear charge still remains the same, and each time an electron is removed the remaining electrons are attracted more strongly. However, a graph of ionisation energy against the number of electrons removed shows the increase as electrons are removed from the second shell is not regular (Figure 1.18). The sub-shells that the electrons are being removed from influence successive ionisation energies. Figure 1.19 plots the successive ionisation energies for magnesium as electrons are removed from the second shell.

paired 2p electrons

0 1 2 3 4 5 6 7 8 9 10 11 12 Number of electrons (magnesium)

Figure 1.19 The third to the tenth successive ionisation energies for magnesium (electrons being removed from the second shell).

Tip

You may find it easier to use a ruler to line up the points in Figure 1.19 to see the patterns.

The third, fourth and fifth electrons are paired electrons in p sub-shells. Because of spin-pair repulsion, these electrons are relatively easy to remove. The sixth, seventh and eighth electrons are unpaired electrons and require relatively more energy to remove. The ninth and tenth electrons are removed from the 2s sub-shell, which is closer to the nucleus, and relatively more energy is required to remove an electron. Using successive ionisation energies Successive ionisation energies can be used to identify an element, its position in the Periodic Table and its electronic configuration.

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1.4 Ionisation energy

25 000 20 000 15 000 10 000

3 4 5 Number of electrons removed

5 000 7

Successive ionisation energy/kJ mol–1

Worked example

Figure 1.20 Successive ionisation energies of an unknown element.

Figure 1.20 shows the first seven successive ionisation energies for an element. Which group in the Periodic Table is the element in?

Answer The first three electrons are removed relatively easily. There is then a relatively large increase in energy required to remove the fourth electron. This electron is removed from an inner shell. The element has three electrons in its outer shell and is in Group 13.

Worked example

These are the first ten successive ionisation energies for an element: Ionisation energy/kJ mol–1 2nd

3rd

1260

2300

3850

4th

5th

6th

7th

8th

9th

10th

5150

6542

9362

11018

33606

38601

43963

1st

a. Which group in the Periodic Table is the element in? b. Which sub-shells are the first ten electrons removed from? c. If the element is between sodium and argon in the Periodic Table, write the electronic configuration for its outer shell of electrons.

Answer a. First, calculate or estimate the increase in ionisation energy as each electron is removed. The answer is probably obvious. Ionisation energy/kJ mol–1

1st

2nd

3rd

4th

5th

6th

7th

8th

9th

10th

1260

2300

3850

5150

6542

9362

11018

33606

38601

43963

The increase in successive ionisation energies needed to remove the first seven electrons is between 1040 and 2820 kJ mol–1

Removing the eighth electron requires an increase of 22 588 kJ mol–1

This element has seven electrons in its outer shell. It is in Group 17. 21

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1 Atomic structure b. An element with seven electrons in its outer shell will have two electrons in an s subshell and five electrons in a p sub-shell. The first five electrons are removed from p sub-shells; the next two electrons are removed from s sub-shells, the next three electrons are removed from p sub-shells in an inner shell. c. Sodium and argon are in Period 3. The electronic configuration for the outer shell is 3s23p5.

Assignment 1.2: Why do scientists think electrons are arranged in shells and sub-shells?

First ionisation energy/ kJ mol–1

519

900

799

1090

1400

1310

1680

2080

Element

One piece of evidence for this theory came from patterns seen in plots of ionisation energy. These are the first ionisation energies for Period 2:

A1. Plot a graph of first ionisation energy against atomic number, Z. A2. Why is there an overall increase in first ionisation energy across Period 2 from lithium to neon? A3. Why is the first ionisation energy of beryllium higher than that of lithium? A4. Why are there dips in the pattern at boron and oxygen? A5. Why is there an increase in the first ionisation energy between: a. boron and nitrogen b. oxygen and neon? A6. If there was a regular increase in the first ionisation energy across Period 2, what might scientists conclude about the existence of sub-shells? A7. What ionisation energy data is needed to provide evidence for the existence of electron shells? A8. How does lithium’s first ionisation energy help to predict its reactivity?

15. Write an equation to show the first ionisation of a sodium atom. 16. Explain why: a. the first ionisation energy of calcium is higher than the first ionisation energy of potassium b. the third ionisation energy of calcium is much higher than the first ionisation energy of calcium c. the first ionisation energy of potassium is lower than the first ionisation energy of sodium d. the first ionisation energy of aluminium is lower than the first ionisation energy of magnesium. 17. An atom of neon and a magnesium ion (Mg2+) have the same electronic configuration. Why does magnesium have a higher third ionisation energy than the first ionisation of argon? 18. Which successive ionisation of magnesium would produce an ion with the electronic configuration 1s22s22p6? 19. Which successive ionisation of fluorine would produce an ion with the electronic configuration 1s22s22p1?

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Learning outcomes

➜➜First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. ➜➜First ionisation energies increase down a group. ➜➜First ionisation energies increase overall across a period. ➜➜Ionisation energies depend on the attraction between the nucleus and the outer electron. ➜➜Ionisation energies depend on the nuclear charge, the atomic/ionic radius, the shielding by inner shells and sub-shells and spin-pair repulsion. ➜➜Successive ionisation energies can be used to determine which group in the Periodic Table the element is in.

chapter overview

Key ideas

Try copying this mini mind map and expanding upon it. Use your notes from other chapters to help you explore how the essential ideas, theories and principles can be linked further together.

Atomic radius Ionic radius

Periodic Table

Protons

Neutrons

Electrons

Trends in ionisation energy

Electron shells and sub-shells s, p, d, f orbitals Energy levels Principal quantum number, n

Atomic number, Z Mass number, A

Atoms and forces

Isotopes

Ionisation energy

Electronic configuration

learning outcomes Now that you have finished this chapter you should be able to: • • • • • • • •

describe atomic structure in terms of protons, neutrons and electrons understand how mass and charge are distributed in an atom define atomic number and mass number describe variations in atomic radius and ionic radius down a group and across a period define the term ‘isotope’ describe chemical and physical properties of isotopes of an element describe how electrons fill shells and sub-shells write full and shorthand electronic configurations for elements up to krypton in the Periodic Table 23

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1 Atomic structure • • • • •

understand and use the ‘electrons in boxes’ notation write equations to represent first, second and subsequent ionisation energies identify trends in first ionisation energy down a group and across a period explain successive ionisation energies use successive ionisation energies to deduce electronic configurations and an element’s position in the Periodic Table.

chapter review

1. How is mass distributed in an atom? 2. How does atomic number differ from mass number for an element? 3. Bromine has two isotopes. What is meant by ‘isotopes of an element’? 4. How does atomic radius vary: a. down a group b. across a period? 5. For Periods 2 and 3, explain how ionic radius depends on whether positive or negative ions are formed. 6. Write the full electronic configuration for the following atoms in the ground state: b. phosphorus. a. argon 7. Write the shorthand electronic configuration for the following ions: b. K+ a. P3– 8. Define first ionisation energy and write a general equation to represent it for an atom M. 9. Describe the trends in first ionisation energy: a. down Group 1 b. across Period 3. 10. Look at these first four successive ionisation energies:

Ionisation energy/kJ mol

–1

1st

2nd

3rd

4th

900

1760

14 800

21 000

a. Identify the group to which this element belongs. b. If the element is in Period 2, write its full electronic configuration.

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