Chemistry notes, chapter 4

The Periodic Table of Elements

Textbook Chemistry Form 4 KSSM (DLP), pages 78–107

What this chapter covers

Describe the historical development of the Periodic Table and the basic principle of arrangement of elements (4.1)
Relate proton number and electron arrangement to the group and period of an element (4.2)
Explain the inert nature of Group 18, the trend in its physical properties and its uses (4.3)
Describe the physical and chemical properties of Group 1 and explain why reactivity increases down the group (4.4)
Describe the physical and chemical properties of Group 17 and explain why reactivity decreases down the group (4.5)
Describe the trends in physical properties and the nature of the oxides across Period 3, and the uses of metalloids (4.6)
State the position, special characteristics and uses of transition elements (4.7)

4.1Development of the Periodic Table

The Periodic Table of Elements classifies the known elements in a table according to certain characteristics. Today elements are arranged in ascending order of proton number (1 to 118), and elements with similar chemical properties are placed in the same vertical column.

Scientist	Contribution	Weakness or importance
Antoine Lavoisier	Grouped elements as gases, non-metals, metals and metal oxides	Inaccurate: included light, heat and some compounds as elements
Johann Dobereiner	Triads: three elements with similar properties, where the atomic mass of the middle one is about the average of the other two (Ca, Sr, Ba; Cl, Br, I)	Worked for a few elements only, but linked properties to atomic mass
John Newlands	Law of Octaves: arranged by increasing atomic mass; properties repeat at every eighth element	Held only for the first 17 elements, but showed a periodic pattern
Lothar Meyer	Graph of atomic volume against atomic mass; similar elements sit at equivalent positions on the curve (alkali metals at the peaks)	Confirmed the repeating pattern of properties
Dmitri Mendeleev	Increasing atomic mass, similar elements in the same column; left gaps for undiscovered elements and predicted their properties	Predictions made the table convincing
Henry Moseley	Used X-ray spectra to find that each element has its own proton number; arranged by increasing proton number	Basis of the modern table; predicted Tc, Pm, Hf and Re

To compare Dobereiner and Newlands: Dobereiner put elements in groups of three with similar properties; Newlands arranged all known elements in a row by atomic mass and found repetition at every eighth element.

Mendeleev and Newlands used atomic mass. Only Moseley used proton number, which is what the modern table uses.

Which scientist proposed the Law of Octaves, and what was it?

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John Newlands. When elements are arranged by increasing atomic mass, the properties of the first element repeat at every eighth element.

4.2Arrangement of elements in the table

Groups are the 18 vertical columns. Periods are the 7 horizontal rows. The valence electrons decide the group; the number of shells filled with electrons decides the period.

Groups 1 and 2: group = number of valence electrons.
Groups 13 to 18: group = number of valence electrons + 10.
Period = number of shells filled with electrons.

Element	Proton number	Electron arrangement	Valence electrons	Group	Period
Lithium, Li	3	2.1	1	1	2
Calcium, Ca	20	2.8.8.2	2	2	4
Aluminium, Al	13	2.8.3	3	13	3
Oxygen, O	8	2.6	6	16	2
Bromine, Br	35	2.8.18.7	7	17	4
Krypton, Kr	36	2.8.18.8	8	18	4

Worked example

Chlorine has proton number 17. Electron arrangement 2.8.7, so 7 valence electrons: group 7 + 10 = 17. Three shells are filled: period 3.

Helium (2) has a full duplet, so it is in Group 18 even though it has only 2 valence electrons. Do not put it in Group 2.

State the electron arrangement, group and period of potassium (proton number 19).

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2.8.8.1. One valence electron, so Group 1. Four shells, so Period 4.

4.3Group 18: the noble gases

Helium, neon, argon, krypton, xenon, radon and oganesson are noble (inert) gases. Their valence shells are full (a stable duplet for helium, an octet for the others), so their atoms do not donate, accept or share electrons. They exist as monoatomic gases.

Trend down the group

Element	Atomic radius (nm)	Melting point (°C)	Boiling point (°C)	Density (g cm⁻³)
Helium	0.050	-270	-269	0.00017
Neon	0.070	-248	-246	0.00080
Argon	0.094	-189	-186	0.00170
Krypton	0.109	-156	-152	0.00350
Xenon	0.130	-122	-108	0.00550

Atomic radius, melting point, boiling point and density all increase down the group. A larger atom means stronger attractive forces between atoms, so more heat energy is needed to overcome them.

Element	Use
Helium	Weather balloons; divers' oxygen tanks
Neon	Advertising board lights
Argon	Fills electric bulbs; inert atmosphere for high-temperature welding
Krypton	Camera flashes; lasers for eye retina treatment; lighthouse lamps
Xenon	Anaesthesia
Radon	Treating cancer

Explain why neon does not react with other elements.

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Its valence shell is full (octet, 2.8), a stable electron arrangement, so its atoms do not donate, accept or share electrons.

4.4Group 1: the alkali metals

Lithium, sodium, potassium, rubidium, caesium and francium. Each atom has one valence electron and donates it to form a +1 ion: M → M⁺ + e⁻.

Physical properties

Soft, shiny grey metals with low density; Li, Na and K float on water.
Low melting and boiling points compared with other metals (iron melts at 1 540 °C).
Down the group: atomic radius increases; melting and boiling points decrease (melting point: Li 186 °C, Na 98 °C, K 64 °C) because weaker attraction between larger atoms needs less heat energy.

Chemical properties (same properties, different reactivity)

Reacts with	Product	Equation (lithium)
Water	Alkaline hydroxide solution and hydrogen gas	2Li + 2H₂O → 2LiOH + H₂
Oxygen	White solid metal oxide (dissolves to give an alkaline solution)	4Li + O₂ → 2Li₂O
Chlorine	White solid metal chloride	2Na + Cl₂ → 2NaCl (sodium)

Litmus: the solution from water or from the oxide turns red litmus paper blue.

Reactivity increases down Group 1. The number of filled shells and the atomic size increase, so the valence electron is further from the nucleus, the nuclear attraction on it is weaker, and the electron is donated more easily.

Rubidium and caesium are very reactive with water and oxygen. Francium is radioactive and unstable.

State the three links in an explanation: more shells, so weaker pull on the valence electron, so it is donated more easily. Always mention the nucleus and the valence electron.

Group 1 metals are kept away from water in the lab; only small pieces are used and the teacher demonstrates the reaction with water.

Write the equation for sodium reacting with water.

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2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

4.5Group 17: the halogens

Fluorine, chlorine, bromine, iodine, astatine and tennessine exist as diatomic molecules. Each atom has seven valence electrons and receives one electron to form a -1 ion: X + e⁻ → X⁻.

Element	State	Colour	Melting point (°C)	Boiling point (°C)	Density (g cm⁻³)
Chlorine	Gas	Greenish yellow	-101	-34	0.003
Bromine	Liquid	Reddish brown	-7	59	3.119
Iodine	Solid	Purplish black	114	184	4.950

Down the group the molecules get larger, intermolecular forces get stronger, so melting point, boiling point and density increase and the colour gets darker.

Reacts with	Product	Example
Water	Acidic solution	Cl₂ + H₂O → HCl + HOCl
Metal	Metal halide	2Fe + 3Br₂ → 2FeBr₃
Alkali	Metal halide + metal halate + water	I₂ + 2NaOH → NaI + NaOI + H₂O

Reactivity decreases down Group 17. A bigger atom has its valence shell further from the nucleus, so the nuclear attraction on an incoming electron is weaker and it is harder to attract an electron to fill the shell.

Fluorine is a pale yellow, very reactive, corrosive and poisonous gas that explodes with hydrogen. Astatine is rare, radioactive and unstable. Uses: chlorine in bleach, bromine in fire extinguishers, iodine as a disinfectant.

Do not copy the Group 1 trend. Reactivity increases down Group 1 but decreases down Group 17: metals lose electrons (easier when the electron is far from the nucleus), non-metals gain electrons (harder when the shell is far from the nucleus).

Why is fluorine more reactive than iodine?

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Fluorine atom is smaller, so its valence shell is closer to the nucleus. The nuclear attraction on an incoming electron is stronger, so it receives an electron more easily.

4.6Period 3

Na, Mg, Al, Si, P, S, Cl and Ar all have three shells filled with electrons, so they are in the same period.

Element	Na	Mg	Al	Si	P	S	Cl	Ar
Atomic radius (nm)	0.186	0.160	0.143	0.118	0.110	0.104	0.100	0.094
Electronegativity	0.9	1.2	1.5	1.8	2.1	2.5	3.0	-

Across the period, atomic size decreases: more protons give a higher nuclear charge, which pulls the electrons in more strongly.
Electronegativity increases for the same reason.
Na, Mg, Al are metals; Si is a metalloid; P, S, Cl, Ar are non-metals. Physical state changes from solid to gas.

Oxides across Period 3

Figure: Chemistry Form 4 KSSM (DLP) textbook, Figure 4.14, p. 99.

Type of oxide	From	Behaviour
Basic	Metals (Na, Mg)	Dissolves to give an alkaline solution; reacts with acid to give salt and water
Amphoteric	Aluminium	Reacts with both acids and alkalis to give salt and water
Acidic	Non-metals (Si, P, S, Cl)	Dissolves to give an acidic solution; reacts with alkali to give salt and water

Metalloids (semi-metals) have properties of both metals and non-metals. They are poor conductors at room temperature but good at high temperature, so silicon is used as a semiconductor in microchips for computers and phones.

Because the oxide of aluminium reacts with both HNO₃ and NaOH, it is amphoteric. Magnesium oxide reacts only with acid (basic); silicon(IV) oxide only with alkali (acidic).

Which element in Period 3, apart from argon, has the smallest atomic radius? Why does the radius fall across the period?

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Chlorine (0.100 nm). Across the period the number of protons rises, so the nuclear charge increases and pulls the electrons in the same three shells closer to the nucleus.

4.7Transition elements

Transition elements are the block between Group 2 and Group 13 (columns 3 to 12), for example Cr, Mn, Fe, Co, Ni and Cu. Scandium and zinc are not counted as transition elements because they lack the special characteristics.

Periodic table with the transition elements, columns 3 to 12, shaded yellow — Position of transition elements. (Textbook Figure 4.16, p. 101)

Figure: Chemistry Form 4 KSSM (DLP) textbook, Figure 4.16, p. 101.

All are metals: shiny solids, very hard, high density, high melting and boiling points, malleable and ductile.

Special characteristic	Example
Act as catalysts	Fe in the Haber process
Form coloured ions and compounds	Cu²⁺ blue, Fe²⁺ green, Fe³⁺ brown, Mn²⁺ pink, MnO₄⁻ purple, Cr₂O₇²⁻ orange, Cr³⁺ green
More than one oxidation number	Fe: +2 and +3; Cu: +1 and +2; Mn: +2, +4, +7
Form complex ions	[Cu(NH₃)₄]²⁺, [Fe(CN)₆]³⁻

Catalyst	Industrial process
Iron, Fe	Haber process (ammonia)
Vanadium(V) oxide, V₂O₅	Contact process (sulphuric acid)
Platinum, Pt	Ostwald process (nitric acid)
Nickel, Ni (or Pt)	Hydrogenation of vegetable oils to margarine

Name two special characteristics of transition elements that other metals such as sodium do not show.

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Any two: used as catalysts; form coloured ions or compounds; have more than one oxidation number; form complex ions.