Bridging 2.indd

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1 Unit 5 Atomic structure Unit 5 Atomic structure (P.2) 5.1 What is an element made of? (P.2) All elements are made of atoms. The atoms of different elements are different. copper cut piece of copper 1 cut cut 2 pieces of copper 2 cut cut 4 pieces of copper 4 etc. etc. Fig. 5.1 Cutting a piece of copper into halves 5.2 Symbols for elements (P.2) Chemists use symbols to represent elements. The symbols for many elements are the first letters of their names. When there are several elements beginning with the same letter, a second letter is used. Some symbols are taken from the Latin names of the elements. 5.2 Element Symbol (first letter of the name) Table 5.1 Symbols for some common elements Element Symbol (two letters from the name) Element Symbol (taken from the Latin name) Carbon C Aluminium Al Copper Cu ydrogen Argon Ar Iron Fe Nitrogen N Calcium Ca Potassium K Oxygen O Chlorine Cl Sodium Na Sulphur S Magnesium Mg Silicon Si 1

2 Topic 2 Microscopic World I 5.3 States of elements (P.3) Elements exist in different states at room temperature and pressure. For example, silver and sulphur are solids, bromine and mercury are liquids, while chlorine is a gas. 5.3 (a) Sulphur a solid element (b) Bromine a liquid element (c) Chlorine a gaseous element Fig. 5.2 The states of some common elements at room temperature and pressure Complete the following table. Element Aluminium Calcium ydrogen Sodium Symbol Cl K O S State at room temperature and pressure 2

3 Unit 5 Atomic structure 5.4 ow to classify elements? (P.5) Metals and non-metals We can classify elements in many different ways. We can use the properties of an element to classify it as a metal or a non-metal. Most of the elements are metals. 5.4 Table 5.2 Comparing the general properties of metals and non-metals Property State at room temperature and pressure Melting and boiling points Appearance Electrical conductivity eat conductivity Effect of bending and hammering Metals solids (except mercury) usually high shiny good conductors good conductors can be bent or hammered into shape Non-metals a few solids (e.g. carbon, sulphur); many gases (e.g. nitrogen, oxygen) but only one liquid (bromine) often low usually dull if they are solids non-conductors (except graphite) poor conductors brittle if they are solids Metalloids Silicon has some properties similar to those of metals and some properties similar to those of non-metals. Silicon belongs to a group of elements called metalloids. 3

4 Topic 2 Microscopic World I Table 5.3 Some properties of silicon Property State at room temperature and pressure Melting and boiling points Appearance Electrical conductivity Effect of bending and hammering solid high Silicon (a metalloid) grey and shiny crystals, or brown powder crystalline form conducts electricity, brown powder form does not conduct electricity brittle 1 Study the following list of elements: carbon, mercury, oxygen, silicon, sodium. a) Which of the above elements is / are metal(s)? / b) Which of the above elements is a metalloid? c) Which of the above elements is / are non-metal(s)? / d) Which non-metal is a solid at room temperature and pressure? 4

5 Unit 5 Atomic structure 2 Study the following descriptions of three elements. Classify each as a metal, non-metal or metalloid. Explain your choice in each case. Element X Y Z Description a yellow solid that melts at 119 C; both the solid and liquid forms do not conduct electricity 119 C a shiny solid which can be bent or hammered into shape easily a shiny brittle solid which can conduct electricity 5.5 Basic structure of an atom (P.8) Atoms consist of a nucleus and a cloud of electrons that move around the nucleus. The nucleus itself contains two types of particles: protons and neutrons. Protons, neutrons and electrons are called subatomic particles. 5.5 cloud of electrons proton neutron Fig. 5.3 A model for the structure of an atom nucleus key: proton neutron 5.6 Atomic number (P.9) The atomic number of an element is the number of protons in an atom of that element. An atom has equal numbers of protons and electrons

6 Topic 2 Microscopic World I Each element has a unique atomic number, which is used to identify the element. Table 5.4 Number of subatomic particles in atoms of the 20 simplest elements 20 Atom Symbol Number of protons (atomic number) Number of neutrons Number of electrons ydrogen elium e Lithium Li Beryllium Be Boron B Carbon C Nitrogen N Oxygen O Fluorine F Neon Ne Sodium Na Magnesium Mg Aluminium Al Silicon Si Phosphorus P Sulphur S Chlorine Cl Argon Ar Potassium K Calcium Ca

7 Unit 5 Atomic structure 5.7 Mass number (P.11) The mass number of an atom is the sum of the numbers of protons and neutrons in the atom ow many protons, neutrons and electrons are there in each of the following atoms? a) 7 3Li b) Al c) Ca 2 Complete the following table. Atom Fluorine Boron Phosphorus Atomic number Mass number protons Number of neutrons 9 10 electrons Isotopes (P.12) 5.8 Isotopes are different atoms of an element which have the same number of protons but a different number of neutrons. Although the isotopes of an element have different masses and physical properties, they have the same chemical properties. Most elements have more than one isotope. 7

8 Topic 2 Microscopic World I Table 5.5 Isotopes of some elements Element Name of isotope Symbol Mass number protons Number of neutrons electrons ydrogen Protium Deuterium Tritium Carbon Carbon Carbon Carbon C 1 3 6C 1 4 6C Relative masses of atoms and the carbon-12 scale (P.13) Relative isotopic mass Atoms are so light that it is difficult to weigh them. Ordinary units are unsuitable for measuring them. Scientists chose the carbon-12 ( 12 6C) atom to be the standard atom. One carbon-12 atom contains six protons and six neutrons and has a mass number of 12. Chemists have defined the carbon-12 atom as having a mass of exactly By comparing the mass of an isotope of an element with this standard, we can obtain the relative isotopic mass of that particular isotope of the element C The relative isotopic mass of a particular isotope of an element is the relative mass of one atom of that isotope on the 12 C = scale. 12 C = The relative isotopic mass of an isotope is roughly equal to its mass number. 8

9 Unit 5 Atomic structure Relative atomic mass Most elements have more than one isotope. The different isotopes of each element have different relative isotopic masses. owever, for the element as a whole, there is only one relative atomic mass. We must consider the relative isotopic mass and relative abundance of each isotope of an element in nature when calculating the relative atomic mass of that element. The relative atomic mass of an element is the weighted average relative isotopic mass of all the naturally occurring isotopes of that element on the 12 C = scale. 12 C = Table 5.6 Relative atomic masses of some common elements Element Symbol Relative atomic mass Element Symbol Relative atomic mass Aluminium Al 27.0 Magnesium Mg 24.3 Calcium Ca 40.1 Oxygen O 16.0 Chlorine Cl 35.5 Potassium K 39.1 Copper Cu 63.5 Silver Ag ydrogen 1.0 Sodium Na 23.0 Iron Fe 55.8 Sulphur S

10 Topic 2 Microscopic World I 1 Calculate the relative atomic mass of magnesium. Isotope 24 Mg 25 Mg 26 Mg Relative abundance (%) % Boron consists of two isotopes: 10 5B and 11 5B. The relative atomic mass of boron is Calculate the relative abundance of the isotopes. 10 5B 11 5B An element X has a relative atomic mass of 87.6, yet no atom of X has this relative mass. Explain. X 87.6X 5.10 The arrangement of electrons in atoms (P.16) Electronic arrangement Electrons move around the nucleus in circular orbits called shells electrons move around the nucleus in shells n = 1 n = 2 n = 3 nucleus Fig. 5.4 Electrons move around the nucleus in shells Each shell can only hold a certain number of electrons. The 1st shell can hold a maximum of 2 electrons. The 2nd shell can hold a maximum of 8 electrons. The 3rd shell can hold a maximum of 18 electrons

11 Unit 5 Atomic structure nucleus 1st shell can hold a maximum of 2 electrons 2[2(1) 2 = 2] 2nd shell can hold a maximum of 8 electrons 8[2(2) 2 = 8] 3rd shell can hold a maximum of 18 electrons 18[2(3) 2 = 18] 4th shell can hold a maximum of 32 electrons 32[2(4) 2 = 32] Fig. 5.5 The maximum number of electrons the first four shells can hold The way in which electrons are arranged in an atom is called its electronic arrangement. We can represent the electronic arrangement of an atom by an electron diagram. A sodium atom has 11 electrons. The first 10 electrons fill up the first and second shells while the last electron goes into the third shell. A sodium atom thus has an electronic arrangement of 2,8, ,8,1 Na sodium atom Fig. 5.6 Electron diagram of sodium atom 1 Draw electron diagrams for atoms of the following elements: a) carbon b) nitrogen c) magnesium 11

12 Topic 2 Microscopic World I 2 Work out the electronic arrangements for atoms of three elements (A to C) from the descriptions given below: A C Element A B C Its atomic number is 8. 8 Description Its atom has 1 completely filled shell and 7 electrons in the second shell. 7 Its atom has 20 electrons. 20 Electronic arrangement 3 An atom of element X has three shells. The electron diagram of its atom is shown below: X X O X? Ex X? X X 5.11 Electrons and orbitals (P.20) Suppose we could photograph the electron in a hydrogen atom at any given moment. The electron is moving at a high speed. The electron would occupy different positions if we took photographs at different moments. If we superimposed millions of such photographs, the resulting picture would resemble a cloud composed of a great number of dots. Thus, in the hydrogen atom, we can imagine the electron as an electron cloud. In theory, there is no sharp boundary to the electron cloud. But we can draw a sphere enclosing about 95% of the cloud. Within the region enclosed by the sphere, there is 95% chance of finding the electron. The region in which there is this high probability of finding the electron is called an orbital % 95% 12

13 Unit 5 Atomic structure ow ideas of the atom developed All chemistry depends on one big idea: that everything is made of atoms. But how did chemists find out about atoms? The atomic theory In 1807, John Dalton put forward an atomic theory to explain the structure of matter and certain aspects of chemical reactions. e suggested that all matter was made of tiny particles, which he called atoms. Discovery of electrons and protons Fig. 5.7 John Dalton I n 1897, J.J. Thomson passed high voltage electricity through a gas in a tube at low pressure. e found that a stream of rays, called cathode rays, moved from the negative electrode to the positive electrode. Besides, cathode rays were also deflected strongly towards the positive plate (Fig. 5.9) John Dalton 1897 J.J. Thomson 5.9 to vacuum pump + + cathode rays Fig. 5.8 J.J. Thomson high voltage deflected towards the + plate Fig. 5.9 Cathode rays are deflected towards the positive plate Thomson suggested that the rays were composed of particles carrying negative charges. e called them electrons. Later Thomson discovered positive particles. e named them protons. 13

14 Topic 2 Microscopic World I The nucleus Rutherford s scattering experiment In 1911, E. Rutherford carried out a number of crucial experiments. Thin sheets of metal foil were bombarded with positively charged particles called alpha particles. Most of the alpha particles went straight through the metal foils. Some were deflected. A few of the alpha particles bounced straight back E. Rutherford nucleus gold atoms beam of alpha particles α Fig E. Rutherford Fig Bombardment of gold foil with alpha particles α Rutherford concluded that most of the atom was empty space with a small, positively charged nucleus at its centre. In 1932, J. Chadwick discovered uncharged particles, which he called neutrons. The proposals of Rutherford and Chadwick led to a model of the atom which was composed of protons, neutrons and electrons. Protons and neutrons were packed closely together in the nucleus. Electrons moved around the nucleus at a considerable distance. Questions 1 Dalton thought that atoms were solid balls. a) Which piece of evidence disproved Dalton s idea? b) ow did this evidence prove that atoms are not solid balls? 2 Why were neutrons more difficult to discover than electrons and protons? 1932 J. Chadwick 1 a) b) 2 14

15 Unit 5 Atomic structure m n e p n s p a n m n i r i m r a m r a s o e a e d o 15

16 Topic 2 Microscopic World I 1 All elements are made of. 2 Chemists use to represent elements. 3 At room temperature and pressure, elements exist in different states (, or ). 4 Elements can be classified into three main groups, and. 5 An atom consists of three types of subatomic particles:, and. The contains protons and neutrons. move around the nucleus in shells. 6 Atomic number of an element = number of in an atom of that element = number of in a atom of that element 7 = number of protons + number of neutrons symbol of an atom 16

17 Unit 5 Atomic structure 8 are different atoms of an element which have the same number of protons but a different number of neutrons. 9 The of a particular isotope of an element is the relative mass of one atom of that isotope on the 12 C = scale. 10 The of an element is the weighted average relative isotopic mass of all the naturally occurring isotopes of that element on the 12 C = scale. 11 The way in which electrons are arranged in an atom is called its. nucleus 1st shell can hold a maximum of electrons 2nd shell can hold a maximum of electrons 3rd shell can hold a maximum of electrons 4th shell can hold a maximum of electrons 12 An is the region in which there is a high probability of finding an electron. 17

18 Topic 2 Microscopic World I Unit 6 The periodic table (P.31) 6.2 The periodic table (P.31) Chemists group elements with similar properties together. This gives rise to the periodic table. In the periodic table, all the elements are arranged in order of increasing atomic number Period 1 Period 2 Group I I Li Group II II Be atomic number name of element symbol Group III III B Group IV IV C Group V V Group VI VI N O Group VII VII F Group 0 0 e Ne Period 3 Na Mg transition metals Al Si P S Cl Ar Period 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Period 5 Period 6 Period 7 A vertical column is called a group A horizontal row is called a period Fig. 6.1 A simplified periodic table showing the first 36 elements 36 Groups the vertical columns in the periodic table The vertical columns in the periodic table are called groups. Groups are numbered from I to VII, followed by Group 0. I VII 0 Table 6.1 Electronic arrangements of atoms of some elements in some groups Group I I Group II II Group VII VII Group 0 0 e (2) Li (2,1) Be (2,2) F (2,7) Ne (2,8) Na (2,8,1) Mg (2,8,2) Cl (2,8,7) Ar (2,8,8) K (2,8,8,1) Ca (2,8,8,2) Br (2,8,18,7) Kr (2,8,18,8) 18

19 Unit 6 The periodic table The group number of an element is equal to the number of outermost shell electrons in its atom. For example, every member of Group I has 1 outermost shell electron in its atom. Atoms with the same number of outermost shell electrons react in a similar way. Elements in the same group have the same number of outermost shell electrons in their atoms, therefore they have similar chemical properties. I 1 The chemical properties of an element depend on the number of outermost shell electrons in its atom. Periods the horizontal rows in the periodic table The horizontal rows in the periodic table are called periods. Table 6.2 Electronic arrangements of atoms of elements in periods 2 and 3 Period 2 Electronic arrangement of atom Period 3 Electronic arrangement of atom Li Be B C N O F Ne 2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8 Na Mg Al Si P S Cl Ar 2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8 Atoms of elements in the same period have the same number of occupied electron shells. The atom of each element in Period 1 has one occupied electron shell. 19

20 Topic 2 Microscopic World I istorical development of the periodic table Tasks You are going to search and present information on the historical development of the periodic table. 1 Your teacher will divide the class into groups. Each group should give a 4-minute presentation. Support your presentation with various aids, such as timelines, diagrams, etc. 2 Write a short report of not more than 300 words to summarize your findings. ints for the search 1 Which chemist arranged elements into groups he called triads? Why did he do that? 2 What is the Law of octaves? Which chemist discovered the Law of octaves? 3 What were the major findings of the Russian chemist Mendeléev? Reference websites 1 Website of the Chemistry Department of the Michigan Technological University Periodic Table Click on Periodic Table to display the relevant articles. 2 An online encyclopedia 2 riodicarrangementsoftheelements.asp 3 Website of the Woodrow Wilson National Fellowship Foundation (an organization that seeks to sponsor academic excellence) Website of Chemsoc, a website providing interesting features and useful services for the chemistry community 4 Chemsoc 20

21 Unit 6 The periodic table 6.3 Patterns across the periodic table (P.36) Across a period, the elements change from metals through metalloids to non-metals. Notice that the reactivity of the elements also changes across a period. Apart from the noble gases, the most reactive elements are near the edges of the periodic table and the least reactive ones are in the centre. 6.3 Table 6.3 Some properties of the elements in Period 3 Element State at room temperature and pressure Melting point ( C) C Boiling point ( C) C Electrical conductivity Type of element Reactivity Group I I Group II II Sodium Magnesium Aluminium solid Silicon Phosphorus Sulphur Chlorine gas Argon reactive good metals moderately reactive moderate metalloid very unreactive moderately reactive Group III III poor non-metals Group IV IV very reactive Group V V Group VI VI extremely unreactive Group VII VII Group 0 0 Period 1 Period 2 Li Be B C N O F e Ne Period 3 Na Mg transition metals Al Si P S Cl Ar Period 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Period 5 Period 6 Period 7 key: metal metalloid non-metal Fig. 6.2 Different blocks of elements in the periodic table 21

22 Topic 2 Microscopic World I 1 An atom of element X has the following structure: X X a) To which period and group of the periodic table does element X belong? X b) What is the atomic number of element X? X c) Refer to a complete periodic table, name element X. X d) Name another element you would expect to have similar chemical properties. X 2 The following is a part of the periodic table. (The letters are NOT the symbols of the elements.) Period 2 Period 3 Period 4 Group I I a b c Group II II Group III III Group IV IV Group V V Group VI VI Group VII VII Group 0 0 d e f g h i j a) Across a period, the elements show a gradual change in some physical properties. Suggest ONE such property. b) ow many outermost shell electrons are there in an atom of g? g c) ow many occupied electron shells are there in an atom of h? h d) Classify each element into a metal, metalloid or non-metal. 22

23 Unit 6 The periodic table 6.4 Group I elements alkali metals (P.38) The six elements in Group I are lithium, sodium, potassium, rubidium, caesium and francium. These elements react with water to form alkalis. ence they are called the alkali metals. The melting point and boiling point of the elements decrease as we move down the group. 6.4 I I Table 6.4 Some physical properties of Group I elements I Element Lithium Sodium Potassium Rubidium Caesium Francium State at room temperature and pressure solid Melting point ( C) C Boiling point ( C) C Density (g cm 3 ) g cm Similarities of Group I elements 1 They all have relatively low melting and boiling points when compared with other metals. 2 They are all soft and can be cut with a knife. 3 They all have low densities lithium, sodium and potassium float on water. 4 They are all reactive metals and must be stored in paraffin oil to prevent them from reacting with the air. 5 They all react vigorously with water to give hydrogen gas and an alkaline solution. 6 They all react with non-metals to form compounds called salts. I 1 I

24 Topic 2 Microscopic World I Differences in reactivity of Group I elements Elements in the same group of the periodic table have similar chemical properties. owever, there is a gradual change in the reactivity of the elements as we move down a group. Group I elements are all very reactive. The reactivity of these elements increases as we move down the group. I I 6.5 Group II elements alkaline earth metals (P.41) The six elements in Group II are beryllium, magnesium, calcium, strontium, barium and radium. These elements are found on the Earth and react with water to form alkalis. ence they are called the alkaline earth metals. 6.5 II II Table 6.5 Some physical properties of the Group II elements II Element Beryllium Magnesium Calcium Strontium Barium Radium State at room temperature and pressure solid Melting point ( C) C Boiling point ( C) C Density (g cm 3 ) g cm Similarities of Group II elements 1 They all have relatively low melting and boiling points when compared with other metals (except Group I metals). 2 They all have low densities. 3 They are all reactive metals and react readily with dilute hydrochloric acid to give hydrogen gas. II 1 I II

25 Unit 6 The periodic table 4 They all react with non-metals to form compounds called salts. Differences in reactivity of Group II elements Group II elements are less reactive than Group I elements. The reactivity increases as we move down the group. 4 II II I Group I I Group II II Li Be melting and boiling points decreasing reactivity increasing Na K Rb Cs Mg Ca Sr Ba Fr Ra melting and boiling points increasing reactivity decreasing Fig. 6.3 A summary of trends of some physical properties and reactivity of Groups I and II elements I II portion o t e periodic table is s own below. Period 1 Period 2 Period 3 Period 4 Group I I Li Be Na Mg K Group II II Ca Group III III B Al Group IV IV C N O F Si Group V V P Group VI VI S Group VII VII Cl Group 0 0 e Ne Ar 25

26 Topic 2 Microscopic World I a) Describe the trend in reactivity down Group I when the elements react with water. I b) Explain why sodium is stored in paraffin oil. c) In which way are the electronic arrangements of atoms of magnesium and calcium i) similar to each other? ii) different from each other? d) Potassium and magnesium are added separately to cold water in troughs. State TWO differences in the observations you expect. 6.6 Group VII elements halogens (P.44) Group VII of the periodic table consists of the non-metals of fluorine, chlorine, bromine, iodine and astatine. These elements react with most metals to form salts. ence they are called the halogens (which mean salt formers). There is a gradual change in state as we move down the group. Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature and pressure. There is also a gradual change in the intensity of colour, from pale yellow to black. 6.6 VII VII Table 6.6 Some physical properties of the first four Group VII elements VII Element State at room temperature and pressure Colour Melting point ( C) C Boiling point ( C) C Fluorine gas pale yellow Chlorine gas greenish yellow Bromine liquid reddish brown Iodine solid black

27 Unit 6 The periodic table Similarities in properties of Group VII elements 1 They are all poisonous and smelly. 2 They are all non-metals. 3 They all react with metals to form compounds called salts. Differences in reactivity of Group VII elements The reactivity of these elements decreases as we move down the group. VII VII 6.7 Group 0 elements noble gases (P.46) The six elements in Group 0 are helium, neon, argon, krypton, xenon and radon. They are called noble gases because they rarely react with other substances. The melting point, boiling point and density of the elements increase as we move down the group Table 6.7 Some physical properties of the Group 0 elements 0 Element elium Neon Argon Krypton Xenon Radon State at room temperature and pressure gas Melting point ( C) C Boiling point ( C) C Density (g cm 3 ) g cm Similarities in properties of Group 0 elements 1 They are all colourless gases at room temperature and pressure. 2 They all have very low melting and boiling points. 3 They are all very unreactive

28 Topic 2 Microscopic World I Stability of Group 0 elements The octet rule suggests that atoms become stable by having eight electrons in their outermost shells (or two electrons in the case of some smaller atoms). We say that these atoms have an octet structure in their outermost shells (or a duplet structure in the case of two electrons). Uses of Group 0 elements Table 6.8 The uses of three Group 0 elements 0 Noble gas elium Neon Argon Use(s) in balloons and airships in advertising signs filling electric light bulbs low density unreactive Reason(s) glows red when an electric current is passed through it does not react with the metal filament in a light bulb Fig. 6.4 elium is used in party balloons Fig. 6.5 Neon is used in advertising signs Fig. 6.6 Argon is used to fill electric light bulbs 28

29 Unit 6 The periodic table The table below lists the atomic numbers of four elements, W, X, Y and Z. WXY Z Element Atomic number W 9 X 10 Y 17 Z 20 a) Suggest an element with an atom that has an octet structure in the outermost shell. b) Two of the elements show similar chemical properties. i) Identify the elements. ii) Explain why the two elements show similar chemical properties. iii) Suggest ONE reaction in which both elements behave similarly. 6.8 Predicting the chemical properties of unfamiliar elements (P.50) We can use the periodic table to predict chemical properties of unfamiliar elements as well. For example, sodium and potassium are Group I elements. Both of them react with water vigorously. Rubidium belongs to the same group. Therefore we expect it to react with water vigorously as well. 6.8 I 29

30 Topic 2 Microscopic World I 1 Mendeléev knew that silicon tetrachloride (SiCl 4 ) existed. Using his periodic table, he correctly predicted the existence of ekasilicon, an element just below silicon in the periodic table. SiCl 4 Predict the chemical formula of the compound formed between ekasilicon and chlorine. 2 A, B and C are three different elements belonging to the same group. The electronic arrangements of their atoms are as follows: AB C Element Electronic arrangement of atom A 2,8,2 B 2,8,8,2 C p,q,18,8,r a) Name the group of elements to which they belong. b) What are the numerical values for p, q and r in the electronic arrangement of an atom of C? C pq r c) Predict the state of element C at room temperature and pressure. C d) Predict which of the elements would react most vigorously with water. Explain your answer. 30

31 Unit 6 The periodic table 6.9 From atoms to ions (P.52) Atoms can obtain the stable electronic arrangements of atoms of noble gases by gaining or losing electrons. Positive ions cations 6.9 A simple ion forms when an atom either loses or gains one or more electrons. It is either positively or negatively charged. When a sodium atom loses an electron, a sodium ion forms. The ion has 11 protons but 10 electrons only. It has one more positive charge than negative charge. Therefore the sodium ion is positively charged. It is a positive ion (i.e. a cation) p 12n 11p 12n + e sodium atom sodium ion Fig. 6.7 A sodium atom loses 1 electron to form a sodium ion with 1 positive charge A sodium ion carries 1 positive charge and is represented by the symbol Na +. The + sign means 1 positive charge. Na + + When an atom of an element loses one or more electrons, it forms a positive ion. A positive ion is called a cation. An atom of a metal can obtain the stable electronic arrangement of an atom of a noble gas by losing one or more electrons. 31

32 Topic 2 Microscopic World I Negative ions anions An atom of a non-metal can obtain the stable electronic arrangement of an atom of a noble gas by gaining one or more electrons. When a chlorine atom gains an electron, a chloride ion forms. The ion has 17 protons but 18 electrons. It has one more negative charge than positive charge. Therefore the chloride ion is a negative ion (i.e. an anion). A chloride ion carries 1 negative charge and is represented by the symbol Cl. The sign means 1 negative charge Cl 17p 18n + e 17p 18n chlorine atom chloride ion Fig. 6.8 A chlorine atom gains 1 electron to form a chloride ion with 1 negative charge When an atom of an element gains one or more electrons, it forms a negative ion. A negative ion is called an anion. An atom of a non-metal can obtain the stable electronic arrangement of an atom of a noble gas by gaining one or more electrons Predicting the charge on an ion (P.55) Atoms of Group I elements lose one electron so as to obtain the electronic arrangements of atoms of noble gases. Ions with 1 positive charge are formed. For atoms of non-metals in Group V, VI and VII, they gain 8 group number electrons in order to obtain the stable electronic arrangements of atoms of noble gases. For example, atoms of Group VI elements form ions with 2 (8 6) negative charges I VVI VII 8 VI

33 Unit 6 The periodic table Positive charge(s) on an ion formed from the atom of a metal = group number of the metal Negative charge(s) on an ion formed from the atom of a non-metal = 8 group number of the non-metal = = 8 1 Consider the element magnesium. a) Draw an electron diagram of a magnesium atom. b) ow can a magnesium atom obtain the electronic arrangement of an atom of a noble gas? c) Draw an electron diagram of a magnesium ion. d) Suggest the charge on a magnesium ion. 2 Consider the following ions: At, Ba 2+, Cs +, Ga 3+, P 3, Se 2 For each ion, state in which group of the periodic table you would expect to find the element which forms the ion. 3 X is an element. It can form a cation X + with an electronic arrangement of 2,8,8. To which period of the periodic table does X belong? Explain your answer. X2,8,8 X + X 33

34 Topic 2 Microscopic World I g p r a m a e m h n g o r d c a 34

35 Unit 6 The periodic table 1 In the periodic table, all the elements are arranged in order of increasing. 2 The vertical columns in the periodic table are called, which are numbered from to, followed by Group (or Group ). Group number of an element = number of in an atom of the element 3 The horizontal rows in the periodic table are called. Period number of an element = number of in an atom of the element 4 Across a period in the periodic table, the elements change from metals through to. 5 a) Elements in the same group have the same number of in their atoms and thus they have similar. b) There is usually a gradual change in the of elements as we move down a group. 6 Group I elements alkali metals a) They all have relatively melting and boiling points when compared with other metals. b) They are all and can be cut with a knife. c) They all have lithium, sodium and potassium float on water. d) They are all metals and must be stored in to prevent them from reacting with air. 35

36 Topic 2 Microscopic World I e) They all react vigorously with water to give gas and an solution. f) They all react with non-metals to form compounds called. g) The reactivity of these elements as we move down the group. 7 Group II elements alkaline earth metals a) They all have relatively melting and boiling points when compared with other metals (except Group I metals). b) They all have densities. c) They are all metals and react readily with to give hydrogen gas. d) They all react with to form compounds called salts. e) Group II elements are reactive than Group I elements. The reactivity as we move down the group. 8 Group VII elements halogens a) They are all and. b) They are all. c) They all react with to form compounds called salts. d) The reactivity of these elements as we move down the group. 36

37 Unit 6 The periodic table 9 The following diagram shows the trends of some physical properties and reactivity of Groups I, II and VII elements. Group I Group II Group VII Li Be F melting and boiling points reactivity Na K Rb Mg Ca Sr Cl Br I melting and boiling points reactivity Cs Ba Fr Ra melting and boiling points reactivity 10 Group 0 elements noble gases a) They are all gases at room temperature and pressure. b) They all have very melting and boiling points. c) They are all very. 11 The suggests that atoms become stable by having eight electrons (an octet structure) in their outermost shells (or two electrons, a duplet structure, in the case of some smaller atoms). 12 Atoms can obtain the stable electronic arrangements of atoms of noble gases by gaining or losing electrons. loses electron(s) atom of positive ion (or ) gains electron(s) atom of negative ion (or ) 37

38 Topic 2 Microscopic World I 13 a) Positive charge(s) on an ion formed from the atom of a metal = b) Negative charge(s) on an ion formed from the atom of a non-metal = 38

39 Unit 7 Ionic and metallic bonds Unit 7 Ionic and metallic bonds (P.68) 7.1 Conductors, electrolytes and non-conductors (P.68) Conductors These are substances which conduct electricity but are not chemically changed during conduction. For example, metals are conductors. Electrolytes These are compounds which conduct electricity in molten state or aqueous solution. They are decomposed by electricity during conduction. Compounds made up of metals and non-metals are electrolytes Examples conductors copper magnesium iron substances electrolytes sodium chloride (made up of sodium and chlorine) lead(ii) bromide (made up of lead and bromine) potassium iodide (made up of potassium and iodine) non-conductors carbon (diamond) chlorine sulphur non-metals compounds distilled water (made up of hydrogen and oxygen) ethanol (made up of carbon, hydrogen and oxygen) sugar (made up of carbon, hydrogen and oxygen) Fig. 7.1 Classification of substances according to how they conduct electricity 39

40 Topic 2 Microscopic World I Non-conductors These are substances which do not conduct electricity in solid, molten state or aqueous solution. All non-metals (except graphite) are non-conductors. Compounds made up of nonmetals are also non-conductors. 7.2 Evidence of ions from electrolysis of molten lead(ii) bromide (P.69) When we pass electricity through the molten lead(ii) bromide, a reddish brown gas (bromine) is formed at the positive electrode. A white shiny solid (lead) is formed at the negative electrode. Lead(II) bromide is decomposed into lead and bromine by electricity. To explain the observations, we assume that lead(ii) bromide is made up of positive lead(ii) ions (Pb 2+ ) and negative bromide ions (Br ). 7.2 (II) (II) (II) (II) (II) Pb 2+ Br Bromide ions carrying negative charges move towards the positive electrode. Lead(II) ions carrying positive charges move towards the negative electrode. (II) + Br Br Pb 2+ Pb 2+ Br Br At the positive electrode Each bromide ion gives up one electron to the electrode and becomes a bromine atom. Bromine atoms then join in pairs to form bromine molecules. bromide ions electrons bromine atoms bromine molecules + e Br Pb 2+ Br e Br Br Pb 2+ At the negative electrode Each lead(ii) ion receives two electrons from the electrode and becomes a lead atom. (II) lead(ii) ions + electrons lead atoms (II) + + Br Br Br Br e Pb 2+ e Pb 2+ Fig. 7.2 Explaining what happens during the electrolysis of molten lead(ii) bromide (II) 40

41 Unit 7 Ionic and metallic bonds In solid state, ions in the compound are held together by strong attraction. They are not free to move. ence solid lead(ii) bromide does not conduct electricity. When lead(ii) bromide becomes molten, the lead(ii) ions and bromide ions become mobile. 7.3 Chemical bonds (P.71) We have learnt that elements combine to form compounds. The particles in these compounds are held together by chemical bonds. A chemical bond is a force that holds the particles together. 7.4 Ionic bonds (P.71) Ionic bond in sodium chloride When a sodium atom loses one electron, it forms a sodium ion with 1 positive charge. (II) (II) (II) Na Na + e sodium atom sodium ion Fig. 7.3 Formation of a sodium ion When a chlorine atom gains one electron, it forms a chloride ion with 1 negative charge. Cl + e Cl chlorine atom chloride ion Fig. 7.4 Formation of a chloride ion 41

42 Topic 2 Microscopic World I When sodium and chlorine react, the electron released by the sodium atom is accepted by the chlorine atom. The compound sodium chloride is produced. + Na Cl Na Cl sodium atom chlorine atom sodium ion chloride ion Fig. 7.5 Electron transfer during the reaction between sodium and chlorine The positively charged sodium ion is attracted to the negatively charged chloride ion by electrostatic forces. This attraction, which holds the ions together, is a chemical bond called an ionic bond. A compound with such a bond is called an ionic compound. An ionic bond is the strong electrostatic forces of attraction between oppositely charged ions. An ionic bond is formed when one or more electrons are transferred from one atom (or group of atoms) to another. When a metal and a non-metal combine to form an ionic compound, atoms of the metal release electrons while atoms of the non-metal gain electrons. Ionic bond in magnesium fluoride A magnesium atom tends to lose two electrons to obtain the electronic arrangement of a stable neon atom. The magnesium ion is represented by the symbol Mg 2+. Mg 2+ 42

43 Unit 7 Ionic and metallic bonds 2+ Mg Mg + 2 e magnesium atom magnesium ion Fig. 7.6 Formation of a magnesium ion A fluorine atom tends to gain one electron to obtain the electronic arrangement of a stable neon atom. The fluoride ion is represented by the symbol F. F F + e F fluorine atom fluoride ion Fig. 7.7 Formation of a fluoride ion When magnesium and fluorine react, the two electrons released by the magnesium atom are accepted by two fluorine atoms. The two negatively charged fluoride ions are attracted to the positively charged magnesium ion. Ionic bonds are formed between the magnesium and fluoride ions. The compound magnesium fluoride is produced. F F Mg fluorine atom Mg 2+ fluoride ion magnesium atom F fluorine atom magnesium ion Fig. 7.8 Electron transfer during the reaction between magnesium and fluorine F fluoride ion 43

44 Topic 2 Microscopic World I Ionic bond in lithium oxide A lithium atom tends to lose one electron to obtain the electronic arrangement of a stable helium atom. The lithium ion is represented by the symbol Li +. Li + + Li Li + e lithium atom lithium ion Fig. 7.9 Formation of a lithium ion An oxygen atom tends to gain two electrons to obtain the electronic arrangement of a stable neon atom. The oxide ion is represented by the symbol O 2. O 2 2 O + 2 e O oxygen atom oxide ion Fig Formation of an oxide ion When lithium and oxygen react, two lithium atoms are required to release the two electrons needed by the oxygen atom. Two positively charged lithium ions are attracted to the negatively charged oxide ion. Ionic bonds are formed between the lithium and oxide ions. The compound lithium oxide is produced. + Li Li 2 lithium atom O lithium ion + O Li oxygen atom Li oxide ion lithium atom lithium ion Fig Electron transfer during the reaction between lithium and oxygen 44

45 Unit 7 Ionic and metallic bonds 1 Use an electron diagram to show the electron transfer when each of the following pairs of elements react. a) sodium and sulphur b) calcium and nitrogen 2 The following table shows the atomic numbers of four elements. Element Atomic number a b c d a) Which TWO elements would form an ionic compound? b) Draw an electron diagram of the compound formed. 7.5 Compounds containing polyatomic ions (P.75) An ion can also be formed from a group of atoms. This is called a polyatomic ion. 7.5 Table 7.1 Examples of polyatomic ions Name ydroxide Nitrate Carbonate Sulphate Ammonium Chemical formula O NO 3 CO 3 2 SO 4 2 N 4 + Model O O O N O O O C S O O O O O N Ionic compounds may contain positive metal ions bonded to negative polyatomic ions. 45

46 Topic 2 Microscopic World I 7.6 Name of ions (P.76) Names of positive ions If a metal forms only one kind of positive ion, the name of the ion is the same as the metal. For example, potassium (K) forms potassium ion (K + ). 7.6 KK + Table 7.2 Names of some common positive ions With 1 positive charge Chemical formula Li + Na + K + Ag + + N 4 + Cu + Name lithium ion sodium ion potassium ion silver ion hydrogen ion ammonium ion copper(i) ion (I) With 2 positive charges Chemical formula Mg 2+ Ca 2+ Zn 2+ Fe 2+ Cu 2+ Pb 2+ Name magnesium ion calcium ion zinc ion iron(ii) ion (II) copper(ii) ion (II) lead(ii) ion (II) With 3 positive charges Chemical formula Al 3+ Fe 3+ Name aluminium ion iron(iii) ion (III) Some metals can form more than one kind of positive ion. For example, copper can form two kinds of positive ions, one carrying 1 positive charge and one carrying 2 positive charges (Cu + and Cu 2+ ). When naming these ions, write a Roman numeral in brackets after the name of the metal to show the number of positive charges. Thus, we use the name of copper(i) ion for Cu +, and copper(ii) ion for Cu 2+. Cu + Cu 2+ Cu + (I) Cu 2+ (II) 46

47 Unit 7 Ionic and metallic bonds Ion Cu + Cu 2+ Fe 2+ Fe 3+ Table 7.3 Examples of metals that can form more than one kind of positive ion Name of ion copper(i) ion (I) copper(ii) ion (II) iron(ii) ion (II) iron(iii) ion (III) Example of compound copper(i) oxide (I) copper(ii) oxide (II) iron(ii) chloride (II) iron(iii) chloride (III) Names of negative ions Negative ions include all simple non-metal ions (except + ) and most polyatomic ions. + Chemical formula F Cl Br I NO 2 NO 3 O CO 3 SO 4 MnO 4 With 1 negative charge Name fluoride ion chloride ion bromide ion iodide ion nitrite ion nitrate ion hydroxide ion hydrogencarbonate ion hydrogensulphate ion permanganate ion Table 7.4 Names of some common negative ions With 2 negative charges Chemical formula O 2 S 2 SO 3 2 SO 4 2 CO 3 2 Cr 2 O 7 2 Name oxide ion sulphide ion sulphite ion sulphate ion carbonate ion dichromate ion With 3 negative charges Chemical formula N 3 PO 4 3 Name nitride ion phosphate ion 47

48 Topic 2 Microscopic World I Simple negative ions have names ending in -ide. Polyatomic ions containing oxygen have names ending in -ite or -ate. The polyatomic ion with less oxygen is named -ite, and that with more oxygen is named -ate. -ide -ite -ate -ite -ate 7.7 Naming ionic compounds (P.78) When naming an ionic compound in English, name the positive ion first, followed by the negative ion. 7.7 Table 7.5 Names of some ionic compounds Positive ion in the compound lithium ion magnesium ion ammonium ion calcium ion copper(ii) ion (II) iron(iii) ion (III) iron(ii) ion (II) potassium ion Negative ion in the compound oxide ion fluoride ion chloride ion nitrate ion carbonate ion hydroxide ion sulphate ion permanganate ion Name of the compound lithium oxide magnesium fluoride ammonium chloride calcium nitrate copper(ii) carbonate (II) iron(iii) hydroxide (III) iron(ii) sulphate (II) potassium permanganate 7.8 Colours of ionic compounds (P.78) If an ionic compound has colour, the colour may arise from either the negative or positive ion, or even from both ions. Consider the colours of the aqueous solutions of potassium chloride and potassium dichromate. The aqueous solution of potassium chloride is colourless. ence the potassium ions must be colourless. Since the aqueous solution of potassium dichromate is orange in colour, the orange colour must come from the dichromate ions

49 Unit 7 Ionic and metallic bonds Table 7.6 Colours of some ions in aqueous solutions Ion Iron(II) (II) Iron(III) (III) Copper(II) (II) Permanganate Dichromate Chromium(III) (III) Nickel(II) (II) Manganese(II) (II) Chemical formula Fe 2+ Fe 3+ Cu 2+ MnO 4 Cr 2 O 7 2 Cr 3+ Ni 2+ Mn 2+ Colour pale green yellow-brown blue or green purple orange green green very pale pink (or colourless) Colours of gemstones The colours of gemstones are due to the presence of traces of coloured ions. Table 7.7 Coloured ions in gemstones Gemstone Colour Ion present Chemical formula of ion green chromium(iii) (III) Cr 3+ jade green chromium(iii) (III) Cr 3+ emerald Continued on next page 49

50 Topic 2 Microscopic World I Gemstone Colour Ion present Chemical formula of ion purple manganese(iii) (III) Mn 3+ amethyst light green iron(ii) (II) Fe 2+ peridot yellow-brown iron(iii) (III) Fe 3+ topaz greenish blue copper(ii) (II) Cu 2+ turquoise Movement of coloured ions When we place a small crystal of potassium permanganate at the centre of a strip of filter paper moistened with tap water, we can see a purple spot slowly moving towards the positive electrode. This is because negative permanganate ions which are purple in colour move towards the positive electrode. Positive potassium ions move towards the negative electrode. owever, we cannot see the potassium ions because they are colourless. Electricity is passed through a gel containing copper(ii) ions and dichromate ions. An orange colour appears near the positive electrode. This is because negative dichromate ions move towards the positive electrode. (II) 50

51 Unit 7 Ionic and metallic bonds purple spot (permanganate ions) original position of potassium permanganate crystal filter paper moistened with tap water microscope slide + 10 V d.c. power supply 10 V Fig Movement of permanganate ions when electricity is passed through A blue colour appears near the negative electrode. This is because positive copper(ii) ions move towards the negative electrode. (II) 24 V d.c. power supply 24 V + carbon electrodes dilute sulphuric acid a gel containing copper(ii) ions and dichromate ions (II) Fig Movement of ions when electricity is passed through a gel containing copper(ii) ions and dichromate ions (II) 51

52 Topic 2 Microscopic World I 1 Topaz is yellow-brown in colour. Suggest the ion responsible for the colour. 2 A student used the following set-up to study the movement of ions. microscope slide filter paper moistened with tap water A B C + d.c. power supply The student placed a drop of copper(ii) sulphate solution at A and a drop of orange solution at C. The two solutions would not react. (II) A C a) The orange colour of the solution at C is due to the anion present. Name the ion responsible for the colour. C b) Electricity was passed through for some time. i) What would be the colour change at A? Explain your answer. A ii) What would be the colour change at B? Explain your answer. B 7.9 Chemical formulae of ionic compounds (P.83) A chemical formula is a way of representing a chemical substance using symbols and figures. The chemical formula of an ionic compound shows: the types of ions present; and the ratio of one type of ion to the other

53 Unit 7 Ionic and metallic bonds Writing chemical formulae of ionic compounds We write the chemical formula of an ionic compound by combining the symbols of its positive and negative ions. Table 7.8 Steps for working out the chemical formulae of ionic compounds Step 1 Write down the symbols of ions in the compound. 2 Write down the number of charges of each ion on the top of each symbol. 3 Cross multiply the numbers and the symbols. 4 Combine the symbols and simplify the ratio if necessary. Calcium oxide Copper(II) hydroxide (II) Iron(III) carbonate (III) Ca O Cu O Fe CO Ca O 2 2 Ca O = Ca 2 = O 2 CaO (Simplify the ratio of 2 : 2 to 1 : 1.) Cu O 2 1 Cu O = Cu 1 = (O) 2 Cu(O) 2 (Omit the number of 1 for Cu.) Fe CO Fe CO 3 = Fe 2 = (CO 3 ) 3 Fe 2 (CO 3 ) 3 1 Write down the names of the following compounds: a) MgCl 2 b) Fe 2 O 3 c) Ca(O) 2 2 Work out the chemical formulae of the following compounds: a) sodium sulphate b) copper(ii) chloride c) ammonium carbonate (II) d) lead(ii) hydroxide e) potassium nitrate f) potassium dichromate (II) 3 M is an element in the third period of the periodic table. It forms a hydroxide which has the chemical formula M(O) 3. What is the chemical formula of the sulphate of M? M M(O) 3 53

54 Topic 2 Microscopic World I 7.10 Metallic bonds in metals (P.85) In a piece of metal, the outermost shell electrons of each atom are not held tightly to the nucleus. Instead, they are free to move randomly in the piece of metal. We can regard the mobile electrons as a sea of electrons. Thus, a piece of metal consists of positively charged ions surrounded by a sea of electrons. The outermost shell electrons are said to be delocalized as they are not associated with a particular ion and can move around e e e e e e e e e e e e e e Fig A piece of metal consists of positively charged ions surrounded by a sea of electrons The attractive forces between the negatively charged electrons and the positively charged ions hold the particles of a metal together. This type of bonding is found only in metals and is called a metallic bond. A metallic bond is a type of bond in which positive metal ions are held together by a sea of mobile electrons. 54