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Chapter 1 Periodic Table and Atomic Structure

Name: ______________________( ) Class: ______ Date: ____________

CHAPTER MAP & OVERVIEW

M. Heyworth Rex, & J G R Briggs. (2013). All About Chemistry 'O' Level. Malaysia: Pearson Education

South Asia Pte Ltd., pages 166 to 171

Learning Outcomes:

Pupils are expected to:

(a) describe the Periodic Table as an arrangement of the elements in the order of increasing proton (atomic) number

(b) describe how the position of an element in the Periodic Table is related to proton number and electronic structure

(c) describe the relationship between group number and the ionic charge of an element (d) explain the similarities between the elements in the same Group of the Periodic Table in terms of

their electronic structure (e) describe the change from metallic to non-metallic character from left to right across a Period of the

Periodic Table (f) describe the relationship between group number, the number of valence electrons and metallic/non-

metallic character

Period

Across the Period: Metallic to Non-

metallic properties

Group

CHAPTER 1.1 PERIODIC TABLE

Down the Group: Same Number of Valance Electrons

Similar Chemical Properties Similar Reactions

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M. Heyworth Rex, & J G R Briggs. (2013). All About Chemistry 'O' Level. Malaysia: Pearson Education

South Asia Pte Ltd., pages 68 to 78

Learning Outcomes:

Pupils are expected to:

(a) state the relative charges and approximate relative masses of protons, neutrons and an electron (b) describe, with the aid of diagrams, the structure of an atom as containing protons and neutrons

(nucleons) in the nucleus and electrons arranged in shells (energy levels) (c) define proton (atomic) and nucleon (mass) number (d) use and interpret such symbols as 126C (e) define the term isotopes and understand the differences in the isotopes of the same element (f) deduce the numbers of protons, neutrons and electrons in atoms and ions given proton and nucleon

numbers (g) Draw the electronic configuration of atoms and ions

Atomic Model

Atomic Number

and Mass Number

Isotopes

Chemical Symbol

Chapter 1.1.4 Electronic

Structure & Configuration

Nucleus

Electron Proton Neutron

Formation of Ions

Relative Atomic Mass

CHAPTER 1.2 ATOMIC STRUCTURE

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1.1 Periodic Table During the 19th century, several chemists looked for patterns in the properties of elements. The most successful of these approaches was by the Russian chemist Dmitri Mendeleev in 1869. Mendeleev arranged all the known elements in order of their relative atomic masses. He also arranged the elements in horizontal rows so that elements with similar properties were in the same vertical column. Because of the periodic repetition of elements with similar properties, Mendeleev called his arrangement a periodic table.

Mendeleev's periodic table The figure below shows part of Mendeleev's periodic table. Notice that elements with similar properties, such as sodium and potassium, fall in the same vertical column. Which other pairs or trios of similar elements appear in the same vertical column of Mendeleevs table?

Mendeleev had some

brilliant and successful ideas in connection with his periodic table.

He left gaps in his table so that similar elements were in the same vertical group. Three of these gaps are shown as asterisks in the figure above.

He predicted the properties of the missing elements from the properties of elements above and below them in his table. Within 15 years of his predictions, the missing elements had been discovered. They were called scandium, gallium and germanium. Their properties were very similar to Mendeleev's predictions.

The success of Mendeleev's predictions showed that his ideas were probably correct. His periodic table was quickly accepted as an important summary of the properties of elements. Mendeleev was the first chemist to successfully arrange the elements into a pattern linking their properties arc relative atomic masses

In the periodic table

The vertical columns of similar elements are called groups.

The horizontal rows of

elements are called periods.

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The modern periodic table The modern periodic table is based on Mendeleev's. It shows all the known elements numbered along each period, starting with period 1, then period 2, etc. The number given to each element is called its atomic number. In the periodic table, the elements are arranged in order of increasing proton (atomic) number, and are classified according to Groups and Periods. Group The groups in the Periodic Table are numbered from I to VII and then Group 0. Some of these groups have names:

Group number Group I Alkali metals II Alkaline earth metals

VII Halogens 0 Noble gases

Elements between Group II and III are known as transition metals or transition elements. Electronic Structure Down each group, the number of valence electrons is the same for each element and is equal to the group number.

Example: Group I Elements

Element Electronic configuration Li 2.1 Na 2.8.1 K 2.8.8.1

Group I elements are very reactive. Since elements with similar electronic configurations have similar chemical properties, elements in the same group have similar chemical properties and will undergo the same type of chemical reactions. Charges on ions Charges on the ions formed are related to the group number and number of valence electrons. Elements on the left side of the Periodic Table lose their valence electrons to form cations with charges corresponding to their group number. Elements on the right side of the Periodic Table gain electrons to form anions. The charges on the anions corresponding to the number of electrons gained to fill their valence shells with eight electrons.

Element Na Mg Al Si P S Cl Ar Group number I II III IV V VI VII 0 Formula of ion Na+ Mg2+ Al3+ - P3- S2- Cl- - Period Each period is numbered, 1, 2, 3, etc.

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Elements in the 1st period will only have their 1st shell fully/partially occupied with electrons. Elements in the 2nd period will have their 1st shell fully occupied with electrons, and their 2nd shell fully/partially occupied with electrons.

Element Proton number

Number of electrons in Electronic configuration

Period Group 1st shell 2nd shell 3rd shell 4th shell

H 1 1 1 1 - He 2 2 2 1 0 Li 3 2 1 2.1 2 I Be 4 2 2 2.2 2 II B 5 2 3 2.3 2 III C 6 2 4 2.4 2 IV N 7 2 5 2.5 2 V O 8 2 6 2.6 2 VI F 9 2 7 2.7 2 VII

Ne 10 2 8 2.8 2 0 Na 11 2 8 1 2.8.1 3 I Mg 12 2 8 2 2.8.2 3 II Al 13 2 8 3 2.8.3 3 III Si 14 2 8 4 2.8.4 3 IV P 15 2 8 5 2.8.5 3 V S 16 2 8 6 2.8.6 3 VI Cl 17 2 8 7 2.8.7 3 VII Ar 18 2 8 8 2.8.8 3 0 K 19 2 8 8 1 2.8.8.1 4 I

Ca 20 2 8 8 2 2.8.8.2 4 II Patterns in the Periodic Table One useful way of classifying elements is as metals and non-metals. Unfortunately, it is not easy to classify some elements in this way. Take, for example, graphite and silicon. These two elements have high melting points and high boiling points (like metals) but they have low densities (like non-metals). They conduct electricity better than non-metals but not as well as metals. Elements with some properties like metals and other properties like non-metals are called metalloids. Because of this difficulty in classifying elements neatly as metals and non-metals, chemists looked for patterns in the properties and reactions of smaller groups of elements.

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Metals and non-metals Across the period, the properties of elements change from metallic to non-metallic.

Generally, elements with small number of electrons in the valence shell (e.g. Group I and II) are metals. Elements with large number of electrons in the valence shell (e.g. Group VII and 0) are non-metals.

The line that divides metals from non-metals runs run diagonally through the Periodic Table. Elements found beside this dividing line (zigzag line) are known as metalloids. Metalloids have some properties of non-metals and metals. Apart from noble gases, the most reactive elements are near the left and right-hand sides of the periodic table. The least reactive elements are in the centre. Sodium and potassium, two very reactive metals, are at the left-hand side. The next most reactive metals, like calcium and magnesium, are in group II, whereas less reactive metals (like iron and copper) are in the centre of the table. Carbon and silicon, unreactive non-metals, are in the centre of the periodic table. Sulfur and oxygen, which are nearer the right-hand edge, are more reactive. Fluorine and chlorine, the most reactive non-metals, are very close to the right-hand edge.

Summary

Group - a vertical set of elements

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Period - a horizontal row of elements

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Modern Chemical Symbols Listed below are the atomic numbers, names, and symbols of the most common elements. The atomic number is used to determine the place of the element in the periodic table; it also has other meaning as you will find out later in the course. Become familiar with the names and symbols of these elements.

Atomic Atomic Number Name Symbol Number Name Symbol

1 hydrogen H 28 nickel Ni 2 helium He 29 copper Cu 3 lithium Li 30 zinc Zn 4 beryllium Be 33 arsenic As 5 boron B 35 bromine Br 6 carbon C 36 krypton Kr 7 nitrogen N 37 rubidium Rb 8 oxygen O 38 strontium Sr 9 fluorine F 47 silver Ag

10 neon Ne 48 cadmium Cd 11 sodium Na 50 tin Sn 12 magnesium Mg 51 antimony Sb 13 aluminum Al 53 iodine 1 14 silicon Si 54 xenon Xe 15 phosphorus P 55 cesium Cs 16 sulfur S 56 barium Ba 17 chlorine Cl 74 tungsten w 18 argon Ar 78 platinum Pt 19 potassium K 79 gold Au 20 calcium Ca 80 mercury Hg 21 scandium Sc 82 lead Pb 22 titanium Ti 83 bismuth Bi 23 vanadium V 86 radon Rn 24 chromium Cr 87 francium Fr 25 manganese Mn 88 radium Ra 26 iron Fe 92 uranium U 27 cobalt Co

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1.2 Atomic Structure What is an atom? An atom is the smallest unit of an element, having the properties of that element. Are there particles smaller than the atom? Atoms are not like solid balls (Figure 1) as proposed by Dalton in 1803.

A century ago, scientists thought that atoms were hard solid particles like incredibly small marbles. In the second part of the 20th century, experiments led to a clearer picture for the structure of atoms.

All atoms are built from just three particles protons, neutrons and electrons.

The centre of an atom is called the nucleus and this contains the protons and neutrons. The nucleus takes up less than 1% of the volume of an atom.

Protons and neutrons have virtually the same mass. The proton and neutron are each assigned a relative mass of one. Protons have a positive charge, but neutrons are neutral.

More than 99% of an atom is empty space occupied by moving electrons. Electrons have a mass about 2 000 times less that of a proton or a neutron.

Electrons have a negative charge. The negative charge on one electron just cancels the positive charge on one proton. Electrons move around very rapidly. They tend to occupy layers or shells at different distances from the nucleus.

The key points about atomic structure are summarised below:

Figure 1: Dalton's model of the atom

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Lets Think Why is the term "relative mass" used in the Table 1 rather than just mass?

Mass of proton, neutron and electron are too small and inconvenient to work with. By using relative mass, we do not have to remember the exact value of the various masses.

1.1.1 The Atomic Model

With reference to the diagram on the left:

(a) The centre of an atom is called the nucleus which contains the protons and neutrons.

(b) The electrons in an atom are arranged in shells (orbits) at different distances from the nucleus. The shell nearest to the nucleus is numbered 1, the second nearest is numbered 2 and so on.

Note: Shells are also called energy levels.

(c) Each shell can hold a certain maximum number of electrons.

i. 1st shell - 2 electrons ii. 2nd shell - 8 electrons

For the 1st 20 elements, the maximum number of electrons that can go into the third shell is 8.

Advanced: For elements after calcium in the 4th period, their third shell can hold up to 18 electrons.

Protons, neutrons and electrons are the building blocks for all atoms. Hydrogen atoms are the simplest atoms. Each hydrogen atom has only one proton and one electron. The next simplest atoms are those of helium with two protons, two electrons and two neutrons. After helium comes lithium, with three protons, three electrons and four neutrons.

Some of the heaviest atoms have large numbers of protons, neutrons and electrons. For example, uranium atoms have 92 protons, 92 electrons and 143 neutrons.

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Notice in all these examples, that an atom always has the same number of protons and electrons. This ensures that the positive charges on the protons balance the negative charges on the electrons and the atom is overall neutral.

If the nucleus of an atom was enlarged to the size of a pea and placed on the top of Nelson's Column, the electrons furthest away would be on the pavement

Atomic number and Mass number

The only atoms with one proton are those of hydrogen. The only atoms with two protons are those of helium. The only atoms with three protons are those of lithium and so on.

This shows that the number of protons in an atom decides which element it is. Because of this, the number of protons in an atom is called its atomic number or proton number. Thus, hydrogen has an atomic number of one, helium has an atomic number of two, lithium three, and so on. Remember also that the order of the element in the periodic table tells you its atomic number. So, chlorine, the seventeenth element in the periodic table with 17 protons and 17 electrons has an atomic number of 17.

The mass of the electrons in an atom is negligible compared to that of the protons and neutrons. Therefore, the mass of an atom can be taken to be the number of protons and neutrons added together. This number is called the mass number or nucleon number of the atom.

Atomic number/ proton number = number of protons Mass number/ nucleon number = number of protons + number of neutrons

For example, aluminium atoms, with 13 protons and 14 neutrons, have an atomic number of 13 and a mass number of 27. Sometimes the symbol Z is used for atomic number and the symbol A for mass number. So, for aluminium, Z = 13 and A = 27.

Chemical Symbol

The figure below shows how the mass number and atomic number are shown with the symbol of an element. In the figure, the element gold (Au) is used as an example.

How many protons, neutrons and electrons are there in a gold atom?

79 protons, 118 neutons and 79 electrons

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Lets Think

An atom can be described as an electrically neutral entity made up of a positively charged nucleus at its centre with negatively charged electrons moving around the nucleus.

(a) Why is the atom electrically neutral? Number of electrons = number of protons. Equal positive and negative charges

(b) Why is the nucleus positively charged? The nucleus contains protons and neutrons. Protons are positively charged while neutrons are electrically neutral.

Isotopes and Relative Atomic Mass

Several elements have relative atomic masses which are whole numbers. For example, the relative atomic mass of carbon is 12.0, that of fluorine is 19.0 and that of sodium is 23.0. This is not surprising, as the mass of an atom depends on the mass of its protons and neutrons, both of which have a relative mass of 1.0. For example, we could calculate the relative mass of fluorine as follows. F!!" atoms have: 9 protons relative mass = 9.0

9 electrons relative mass = 0.0 10 neutrons relative mass = 10.0 Therefore, relative atomic mass of F!!" = 19.0

Unlike fluorine, carbon and sodium, some elements have relative atomic masses that are nowhere near whole numbers.

For example, the relative atomic mass of chlorine is 35.5 and that of copper is 63.5. These unexpected results were explained in 1919 when W.F. Aston built the first mass spectrometer.

Using his mass spectrometer, Aston found that one element could have atoms with different masses.

These atoms of the same element with different masses are called isotopes.

Isotopes are atoms with the same atomic number, but different mass numbers.

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The photo above shows evidence for the two isotopes in neon, neon-20 and neon-22. The trace for neon-20 is much thicker than that for neon-22. What does this tell you about the two isotopes? What does CO represent? Each isotope has a relative atomic mass which is a whole number, but the average relative atomic mass for the mixture of isotopes is not always a whole number.

Chlorine is a good example of an element with isotopes. Naturally occurring chlorine contains two isotopes. Cl!"!" is called chlorine-35 and Cl!"!" is called chlorine-37. Each of these isotopes has 17protons and 17 electrons. Therefore, both isotopes have the same atomic number and the same chemical properties because these are determined by the number of electrons. However, one isotope ( Cl!"!" ) has 18 neutrons and the other ( Cl!"!" ) has 20 neutrons. Therefore, they have different mass numbers, different masses and hence different physical properties because these depend on the masses of atoms and molecules.

The similarities and differences between isotopes of the same element are summarised in the table below.

Isotopes can be divided into two types. One type is radioactive; the other is non-radioactive. Radioactive isotopes give out radiation. This radiation is invisible but it can be detected with special instruments. Radiation is harmful to life, and in large amounts, can kill people. It was radioactive isotopes from the Chernobyl nuclear reactor accident that caused the pollution in 1986. Workers handling radioactive isotopes must take precautions to shield themselves from radiation. Most isotopes in the air and the ground are non-radioactive and they do not produce radiation. Most radioactive isotopes are made artificially. Radioactive isotopes have important uses. The cobalt isotope, 60Co, is used in hospitals to treat cancer patients. The intense radiation from this isotope destroys the cancer cells. Cobalt-60 is also used to sterilise surgical instruments used in hospital operations. The powerful radiation kills germs.

Radioactive isotopes produce heat which can be used to produce electrical energy. Small amounts of radioactive isotopes are used to supply energy in remote places. Radioactive isotopes provide energy for spacecraft exploring the outer planets, such as Jupiter and Saturn.

Some people have irregular heartbeats. They need to have a heart pacemaker implanted inside their chest. This instrument provides a tiny electrical shock to ensure a steady heartbeat. Pacemakers can be powered by radioactive isotopes. Such a pacemaker can work reliably for over 20 years. An ordinary battery would have to be replaced every 10 years.

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Calculating relative atomic masses (Optional)

The relative atomic mass of an element is the average mass of one atom. This can be calculated from the relative masses of its isotopes and their relative proportions.

Look closely at the figure below which shows a mass spectrometer trace for chlorine. The trace shows that chlorine contains two isotopes, with mass numbers of 35 and 37.

What are the relative amounts of isotopes Cl!"!" and Cl!"!" ? If chlorine contained 100% Cl!"!" , its relative atomic mass would be 35. If chlorine contained 100% Cl!"!" its relative atomic mass would be 37. If chlorine contained 50% 35 17 C1 and 50% 37 17 Cl, the relative atomic mass would be:

Now, the previous graphical figure shows that chlorine contains three times as much Cl!"!" as Cl!"!" i.e. the percentages of the two isotopes are 75% to 25%. Therefore the atomic mass of chlorine is given by:

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Electronic Structure & Electronic Configuration

Nitrogen-14 atom has 7 electrons. Two of its electrons will go into the 1st shell; the remaining electrons will go into the 2nd shell (Figure 2). With 7 electrons, nitrogen has the electronic configuration of 2.5 Argon-40 atom has 18 electrons. Two of its electrons will go into the 1st shell, 8 electrons will go into the 2nd shell, and the remaining 8 electrons will go into the 3rd shell (Figure 3). With 18 electrons, argon has the electronic configuration of 2.8.8

Figure 2: Electronic Structure (Full Electronic Configuration) of Nitrogen-14 atom

Figure 3: Electronic Structure (Full Electronic Configuration) of Argon-40 atom

Valence Shell (Outer Shell) The shell which is farthest from the nucleus and occupied by electrons is called the valence shell (outer shell). The electrons in the valence shell are known as valence electrons (outer electrons). In a chemical reaction, only these valence electrons are involved in chemical bonding between atoms. Often, only the valence electrons are drawn in the electronic structure. This is called the outer electronic structure. An example is shown in Figure 4 for the Nitrogen-14 atom. Figure 4

N Ar

N

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Formation of Ions

During chemical reactions, some atoms might lose/gain electron(s). Atom becomes an ion (charged particle) when it gains or loses electron(s). Question: Why does an atom become a charged particle when it gains or loses electron(s)? An atom is electrically neutral because number of electrons = number of protons (equal positive and negative charges). When it gains or loses electron(s), the positive and negative charges are not balanced. Therefore, the atom becomes a charged particle.

(a) Formation of Cations When an atom loses one or more electrons, it becomes a positively charged particle called cation.

Lithium atom (Li)

3 electrons 3 protons

Net charge: 0

Lithium ion (Li+) 2 electrons 3 protons

Net charge: +1 In a lithium atom, there are 3 protons and 3 electrons. In a lithium ion, there are 3 protons and 2 electrons. Therefore, the lithium ion carries an overall positive charge of 1+ and is written as Li+. (b) Formation of Anions When an atom gains one or more electrons, it becomes a negatively charged particle called anion.

Fluorine atom (F)

9 electrons 9 protons

Net charge: 0

Fluoride ion (F-) 10 electrons

9 protons Net charge: -1

In a fluorine atom, there are 9 protons and 9 electrons. In a fluoride ion, there are 9 protons and 10 electrons. Therefore, the fluoride ion carries an overall positive charge of 1- and is written as F-.

Li Li

F F