GCSE Chemistry eGuide Unit 1 PDF

GCSE CHEMISTRY: GUIDANCE eGUIDE// GCSE Chemistry: eGuide Unit 1: Structures, Trends, Chemical Reactions, Quantitative Chemistry and Analysis Content/Specification Section Page 1.1 Atomic Structure 02 1.2 Bonding 15 1.3 Structures...

GCSE CHEMISTRY: GUIDANCE eGUIDE// GCSE Chemistry: eGuide Unit 1: Structures, Trends, Chemical Reactions, Quantitative Chemistry and Analysis Content/Specification Section Page 1.1 Atomic Structure 02 1.2 Bonding 15 1.3 Structures 24 1.4 Nanoparticles 34 1.5 Symbols, formulae and equations 36 1.6 The Periodic Table 46 1.7 Quantitative chemistry 60 1.8 Acids, bases and salts 74 1.9 Chemical analysis 88 1.10 Solubility 101 Glossary terms 106 Test yourself answers 109  GCSE CHEMISTRY: GUIDANCE 1.1 Atomic Structure This section will examine the history and basic structure of the atom with reference to subatomic particles, the size of atoms and their nuclei and relative atomic mass. Mathematical content Students will calculate the numbers of subatomic particles in atoms and ions, work with nanometre (nm), picometre (pm) and femtometre (fm) scale in terms of atoms and their nuclei and calculate relative atomic masses of elements from mass numbers and abundance. Learning outcomes demonstrate knowledge and understanding of how ideas about the atom changed over time, with reference to: – the Plum Pudding model; – Rutherford’s model of a nucleus surrounded by electrons; and – the discovery of the neutron by Chadwick, leading to today’s model of an atom. describe the structure of an atom as a central positively charged nucleus containing protons and neutrons (most of the mass) surrounded by orbiting electrons in shells state the relative charges and approximate relative masses of protons, neutrons and electrons define atomic number as the number of protons in an atom define mass number as the total number of protons and neutrons in an atom demonstrate knowledge and understanding that an atom as a whole has no electrical charge because the number of protons is equal to the number of electrons calculate the number of protons, neutrons and electrons in an atom or an ion and deduce the charge on an ion or determine the number of subatomic particles given the charge write and draw the electronic configuration (structure) of atoms and ions with atomic number 1–20 recall that atoms have a radius of about 0.1 nm (1 × 10-10 m) and that the nucleus is less than 1/10 000 of that of the atom (less than 1 × 10-14 m) define isotopes as atoms of an element with the same atomic number but a different mass number, indicating a different number of neutrons interpret data on the number of protons, neutrons and electrons to identify isotopes of an element calculate the relative atomic mass of elements from the mass number and abundances of its isotopes recall that a compound is two or more elements chemically combined pg 2  GCSE CHEMISTRY: GUIDANCE Atomic Structure Introduction Atoms are the smallest particles of matter which can exist on their own in a stable environment. Atoms consist of subatomic particles (protons, neutrons and electrons). The organisation of these subatomic particles gives atoms mass and defines their chemical properties. History of development of the atom Atoms were first described in 1803 as spheres which could not be divided. The electron was discovered about 100 years later and this led to the plum pudding model of the atom. This described the atom as a positive sphere with electrons embedded in it. This shows a representation of the plum pudding model of the atom with the positive sphere and the electrons embedded in it like the raisins in a plum pudding. Ernest Rutherford revised this model to have the nucleus containing protons at the centre of the atom with the electrons orbiting around it. The mass of the atom was centred in the nucleus but it was James Chadwick who discovered the neutron in 1936. The neutron was more difficult to discover as it was not charged unlike the proton and the electron. Rutherford’s model of the atom with the electrons orbiting a central nucleus. History of the development of theories about atoms: https://www.compoundchem.com/2016/10/13/atomicmodels/ pg 3  GCSE CHEMISTRY: GUIDANCE Atomic Structure Subatomic particles The model of the atom we use today comprises all these developments. The atom is composed of a central nucleus containing protons and neutrons with electrons orbiting the nucleus in shells. The nucleus is positively charged due to the protons. The electrons in the shells are negatively charged. The electrons are held by the attractive power of the nucleus. The table below shows the relative masses, relative charges and location of the subatomic particles. Subatomic Location in the Relative mass Relative charge particle atom proton 1 +1 nucleus 1 shells orbiting electron -1 1840 the nucleus neutron 1 0 nucleus The relative masses of the proton and the neutron are the same but the mass of an electron is so small that it does not affect the mass of an atom or ion formed from the atom. The mass of an atom is centred in the nucleus. Exam Tip Be sure to learn the values in the table as they are often asked. Always write the charges as +1 and -1 and not just + and – as the number is important as well. Atomic and nuclear radius The radius of the nucleus of an atom is at least 1/10000 times smaller than the radius of the atom. The radius of an atom is often measured in nanometres (nm). 1 nm = 1 × 10-9 m. There are 1 000 000 000 nm in 1 metre. Most atoms have a radius around 0.1 nm (1 × 10-10 m) and the radius of most nuclei are smaller than 1 × 10-14 m. For example, the atomic radius of a sodium atom is 0.186 nm. The radius of the nucleus of a sodium atom is approximately 54500 times smaller than the atomic radius. 1.86 × 10-10 Nuclear radius = = 3.41 × 10-15 m. 54500 pg 4  GCSE CHEMISTRY: GUIDANCE Atomic Structure Symbols for atoms and ions When writing the formula of an atom or a simple ion, use the symbols which are given on the Periodic Table. Hydrogen for example is represented by H. H represents a hydrogen atom. H+ represents a hydrogen ion showing its charge (positive simple ions keep the same name as the atom from which they are formed). H- is a hydride ion (simple negative ions change their name to end in ide). Charges can be + or - (meaning 1+ and 1- but should be written just + and -) also 2+ or 2- and 3+ or 3-. It is unusual to get a simple ion with a charge larger than 3. Always write the number (apart from 1) before the + or the -. Some examples of atoms and ions are shown in the table below. Atom/ion name Symbol magnesium atom Mg magnesium ion Mg2+ chlorine atom Cl chloride ion Cl- oxygen atom O oxide ion O2- Don’t worry too much about the size of the charge yet as you will understand this better as you work through this unit. pg 5  GCSE CHEMISTRY: GUIDANCE Atomic Structure Atomic number, mass number and number of subatomic particles The atomic number is the number of protons in the nucleus of an atom. The mass number is the total number of protons and neutrons in the nucleus of an atom. An atom has no electrical charge as the number of protons is equal to the number of electrons. This means the total +1 charges of the protons cancel out the total -1 charges of the electrons. The mass number and atomic number may be used to determine the number of subatomic particles in an atom. Often the atom (or ion) is written with its atomic number and its mass number. For example, 27 13 Al. The top number is the mass number of this atom and the bottom number is the atomic number. This means that this aluminium atom has 13 protons (equal to the atomic number) and 27 protons and neutrons in total so there must be 14 neutrons (27 – 13). As it is an atom, the number of protons is equal to the number of electrons so there are 13 electrons. For simple ions, the charge on the ion is equal to the number of protons minus the number of electrons. This is because the ion is charged when an atom loses or gains electrons. Loss of electrons creates a positive ion and gain of electron creates a negative ion. This means that in a simple ion, the number of protons is not equal to the number of electrons. This can all be remembered in four rules: 1. number of protons = atomic number 2. number of neutrons = mass number – atomic number (Also rearranging this gives: mass number = atomic number + number of neutrons or mass number = number of protons + number of neutrons) 3. in an atom, number of protons = number of electrons 4. in a simple ion, the charge of the ion = number of protons – number of electrons Exam Tip Remember that for an atom and its simple ion, the atomic number is the same and so is the mass number. The only things that change are the number of electrons and the charge. pg 6  GCSE CHEMISTRY: GUIDANCE Atomic Structure WORKED EXAMPLES 1. An atom of fluorine has atomic number 9 and mass number 19. State the number of each subatomic particle in the atom. Answer: Fluorine atom has 9 protons, 10 neutrons and 9 electrons. Explanation: The number of protons = atomic number (Rule 1) so 9 protons The number of neutrons = mass number – atomic number = 19 – 9 = 10 neutrons (Rule 2) The number of electrons = the number of protons as this is an atom so 9 electrons (Rule 3) Exam Tip Make sure you don’t get neutrons and electrons mixed up as this is a common mistake. 2. A simple ion formed from an atom has a charge of 2+ and has 12 protons and 12 neutrons. (a) Identify the atom from which the ion was formed. (b) How many electrons are present in the ion? (c) What is the mass number of the ion? Answer: (a) magnesium (b) 10 (c) 24 Explanation: The atomic number identifies the atom or ion and this particle has 12 protons so the atomic number is 12 (Rule 1). If you look up your Periodic Table the element with atomic number 12 is magnesium. Exam Tip Make sure you can use your Periodic Table properly as students often read the wrong number and in this case, you might think it was carbon (C) as this has a (relative atomic) mass of 12 on the Periodic Table but magnesium (Mg) has an atomic number of 12. The charge on the ion is 2+ and this equals the number of protons minus number of electrons so +2 = 12 – number of electrons so the number of electrons must be 10 (Rule 4). The number of neutrons = mass number – atomic number so 12 = mass number – 12 so mass number = 12 + 12 = 24 (Rule 2). pg 7  GCSE CHEMISTRY: GUIDANCE Atomic Structure 3. Complete the following table. Atomic Mass Number of Number of Number of Atom/ion number number protons neutrons electrons C 6 12 P 15 16 8 16 10 Cl- 35 19 39 18 This is a very common style of question and you must use all the rules to complete it. Answer: Atomic Mass Number of Number of Number of Atom/ion number number protons neutrons electrons C 6 12 6 6 6 P 15 31 15 16 15 O2- 8 16 8 8 10 Cl- 17 35 17 18 18 K+ 19 39 19 20 18 Explanation: In row 1 the number of protons = atomic number = 6; number of neutrons = mass number – atomic number = 12 – 6 = 6; it is an atom as the particle on the left has no charge so number of protons = number of electrons. In row 2 the number of protons = atomic number = 15; the mass number = number of protons + number of neutrons = 15 + 16 = 31; It is an atom of phosphorus so 15 electrons (= to the number of protons). In row 3 the atomic number = number of protons = 8; the number of neutrons = mass number – atomic number = 16 – 8 = 8; the particle is an ion so charge = number of protons – number of electrons so -2 = 8 – number of electrons so there must be 10 electrons. pg 8  GCSE CHEMISTRY: GUIDANCE Atomic Structure In row 4 the particle is an ion formed from chlorine so its atomic number on the Periodic Table is 17; this means that there are 17 protons; number of neutrons = mass number – atomic number = 35 – 17 = 18; The ion has a 1- charge so -1 = 17 – number of electrons so 18 electrons. In row 5 the atomic number = number of protons = 19; neutrons = 39 – 19 = 20; particle is a potassium ion as its atomic number is 19 (check Periodic Table) and the number of electrons is not equal to the number of protons; the charge = number of protons – number of electrons = 19 – 18 = 1 or +1 so it is K+. Test Yourself 1.1.1 1. State the relative charge and relative mass of the following subatomic particles. (a) proton (b) neutron (c) electron 2. Name the following simple particles. (a) Cl (b) K+ (c) N3- (d) Al (e) F- (f) Ca2+ (g) H- 3. Write the formula of the following atoms and ions (including any charge on the ion) which contain the following subatomic particles. State the mass number of these particles. (a) 3 protons 4 neutrons 3 electrons (b) 11 protons 12 neutrons 10 electrons (c) 16 protons 16 neutrons 18 electrons 4. Complete the table below. Atomic Mass Number of Number of Number of Atom/ion number number protons neutrons electrons Si 28 14 Mg2+ 12 24 P3- 15 16 5. The nuclear radius of a potassium atom is 4.07 × 10-15 m. The atomic radius is 56700 time larger. Calculate the atomic radius in nm. Give your answer to 3 significant figures. pg 9  GCSE CHEMISTRY: GUIDANCE Atomic Structure Electronic configuration The electrons in an atom are arranged in shells. The shells are at increasing distance from the nucleus. The first shell is closest to the nucleus and it can contain a maximum of 2 electrons. The second shell is further from the nucleus and it can contain a maximum of 8 electrons. The third shell is further again from the nucleus and it can contain a maximum of 8 electrons. The fourth shell is further again from the nucleus and it can contain a maximum of 18 electrons. Exam Tip The maximum number of electrons can be seen in the pattern of elements in the Periodic Table. There are 2 elements in the first period (first row – hydrogen and helium); there are 8 elements in the second period (lithium to neon) and 8 in the third period (sodium to argon) and 18 in the fourth period (potassium to krypton). Electrons are added to shells closest to the nucleus until these shells are full and then the next shell starts to fill with electrons. Electrons pair up but in shells with 8 electrons they only pair up when there are more than 4 electrons. The electrons in each shell can be represented by a drawing. Circles are drawn to represent the shells at increasing distance from the nucleus. Electrons are represented on the circles as a dot ( ) or a cross (×). Crosses are used when drawing a single atom but dot and crosses are used when showing bonding which will be discussed later. The arrangement of electrons may also be shown in written format such as 2, 8, 5 which would indicate that there are 2 electrons in the first shell, 8 electrons in the second shell and 5 electrons in the third shell. The arrangement of electrons in written form or as a diagram is called the electronic configuration. In a shell with capacity for 8 electrons, there are 4 spaces for the electron pairs in a diagram (at 12, 3, 6 and 9 o’clock on the circle) and electrons only pair when there are more than 4 electrons so 4 electrons would have 1 electron positioned at these 4 points on the circle. pg 10  GCSE CHEMISTRY: GUIDANCE Atomic Structure The electronic configurations of atoms of elements in Period 2 are shown below. lithium beryllium boron carbon nitrogen oxygen fluorine neon × × × × ×× ×× ×× ×× ×× ×× ×× ×× ×× ×× ×× ×× × × × × × × × × × × × × × × × × × × × × ×× ×× 2, 1 2, 2 2, 3 2, 4 2, 5 2, 6 2, 7 2, 8 When atoms form simple ions, they lose or gain electrons to obtain a full outer shell. A full outer shell of electrons is stable. For example, a magnesium atom has an electronic configuration of 2, 8, 2. To obtain a full outer shell, the atom reacts and loses the 2 outer electrons so the ion formed is Mg2+ and has an electronic configuration of 2, 8. The number of protons in the nucleus remain the same so it is still a magnesium particle. A chlorine atom has an electronic configuration of 2, 8, 7. When it reacts a chlorine atom gains one electron so it becomes a chloride ion, Cl- with an electronic configuration of 2, 8, 8. WORKED EXAMPLES 1. Write electronic configurations for the following atoms and ions. (a) lithium atom, Li (b) potassium ion, K+ (c) oxide ion, O2- Answers: (a) 2, 1 (b) 2, 8, 8 (c) 2, 8 Explanation: In (a) lithium has 3 electrons which are arranged as 2, 1. In (b) a potassium atom has 19 electrons arranged 2, 8, 8, 1 so when the atom forms the K+ ion, it loses its outer shell electron and the electronic configuration is 2, 8, 8. In (c) an oxygen atom has an electronic configuration of 2, 6 so it gains 2 electrons to form O2- and the electronic configuration is 2, 8. pg 11  GCSE CHEMISTRY: GUIDANCE Atomic Structure 2. Draw a diagram of an atom of phosphorus (atomic number 15, mass number 31) showing the position and number of all subatomic particles. Answer: × nucleus containing 15 ×× × ×× protons and 16 neutrons ×× × × ×× ×× × Explanation: The atom of phosphorus has 15 protons, 16 neutrons and 15 electrons. The first shell has 2 electrons, the second shell has 8 electrons and the third shell has 5 so its electronic configuration is 2, 8, 5. 3. Complete the table below. Number of Number of Electronic Atom/ion protons electrons configuration Be 7 10 S2- 16 12 2, 8 K+ 19 Answer: Number of Number of Electronic Atom/ion protons electrons configuration Be 4 4 2, 2 N3- 7 10 2, 8 S2- 16 18 2, 8, 8 Mg2+ 12 10 2, 8 K+ 19 18 2, 8, 8 pg 12  GCSE CHEMISTRY: GUIDANCE Atomic Structure Explanation: Be has atomic number 4 so it has 4 protons and 4 electrons as it is an atom. The second row has to be an ion formed from nitrogen as it has atomic number 7 and has 10 electrons so has a charge of 3-. This is a nitride ion and the 10 electrons are arranged 2, 8. The S2- ion must have 2 more electrons than protons so it has 18 electrons which are arranged 2, 8, 8. The fourth row must be a magnesium ion as it has 12 protons and 10 electrons (based on 2+8). So it is Mg2+. The last K+ must have 18 electrons arranged 2, 8, 8 as it has 19 protons so 18 electrons will give it an overall charge of +. Isotopes and relative atomic mass Isotopes are atoms of the same element (same atomic number) with a different number of neutrons (different mass number). Chlorine exists as two isotopes 35Cl and 37Cl. Both are atoms of chlorine with 17 protons and 17 electrons, arranged 2, 8, 7 but 35Cl has 18 neutrons and 37Cl has 20 neutrons. Both atoms react in the same way as they have the same electronic configuration but one atom is heavier than the other due to the extra 2 neutrons. The masses of all atoms are measured relative to the mass of an atom of carbon-12 which is used as the standard. The relative atomic mass is the weighted average of the masses of all the atoms of isotopes of an element relative to the mass of an atom of carbon-12. The relative atomic mass of chlorine is 35.5 and this is the figure given on the Periodic Table as it is a weighted average of the masses of the 35Cl and 37Cl atoms. Chlorine atoms occur naturally as 75% 35Cl and 25 % 37Cl. If we take an average using a sample of 100 atoms, 75 would have a mass of 35 and 25 would have a mass of 37. (75×35)+(25×37) relative atomic mass = = 35.5 100 This type of calculation can be carried out for any element if you are given percentage abundance and the mass numbers of the isotopes. Exam Tip Apart from chlorine, all relative atomic masses on the Periodic Table provided in your exams are given to the nearest whole number. When elements react together, they form compounds. A compound is two or more elements chemically combined. pg 13  GCSE CHEMISTRY: GUIDANCE Atomic Structure Test yourself 1.1.2 1. Identify the following atoms and ions, including any charge on the ion. (a) A has 8 protons and an electronic configuration of 2, 8. (b) B has 10 protons and an electronic configuration of 2, 8. (c) C has 13 protons and an electronic configuration of 2, 8. 2. Draw a labelled diagram of a potassium ion, K+ (atomic number 19, mass number 39). 3. Complete the table below. Ion Atomic number Electronic configuration Ca2+ F- Li+ H- 4. Sulfur (atomic number 16) has three isotopes as shown in the table below. Isotope Percentage abundance 32S 95.0 33S 0.75 34S 4.25 (a) What is meant by the term isotopes? (b) Calculate the number of neutrons in each of the isotopes of sulfur. (c) Calculate the relative atomic mass of sulfur. Give your answer to 1 decimal place. pg 14  GCSE CHEMISTRY: GUIDANCE 1.2 Bonding This section will explore the three types bonding – ionic, covalent and metallic. Mathematical content Students will work with different numbers of outer shell electrons to determine the formula of ionic and molecular covalent compounds. Learning outcomes Ionic bonding demonstrate knowledge and understanding that an ion is a charged particle formed when an atom gains or loses electrons and a molecular ion is a charged particle containing more than one atom define the terms cation and anion and explain, using dot and cross diagrams, how ions are formed and how ionic bonding takes place in simple ionic compounds, limited to elements in Groups 1 (I) and 2 (II) with elements in Groups 6 (VI) and 7 (VII), the ions of which have a noble gas electronic configuration demonstrate knowledge and understanding that: – ionic bonding involves attraction between oppositely charged ions – ionic bonds are strong – substantial energy is required to break ionic bonds recognise that ionic bonding is typical of metal compounds Covalent bonding describe a single covalent bond as a shared pair of electrons explain, using dot and cross diagrams, how covalent bonding occurs in H2, Cl2, HCl, H2O, NH3, CH4 and similar molecules and label lone pairs of electrons draw dot and cross diagrams and indicate the presence of multiple bonds in O2, N2 and CO2 recognise covalent bonding as typical of non-metallic elements and compounds demonstrate knowledge and understanding that a molecule is two or more atoms covalently bonded and that diatomic means that there are two atoms covalently bonded in a molecule demonstrate knowledge and understanding that covalent bonds are strong and that substantial energy is required to break covalent bonds demonstrate knowledge and understanding that a covalent bond may be represented by a line Metallic bonding demonstrate knowledge and understanding that metallic bonding results from the attraction between the positive ions in a regular lattice and the delocalised electrons pg 15  GCSE CHEMISTRY: GUIDANCE Bonding Ionic bonding Ionic bonding is the attraction between oppositely charged ions. Ionic bonding is strong and requires substantial energy to break the attraction between the ions. Ionic bonding is typical of metal compounds. An ion is a charged particle. The general term for a positive ion is a cation; the general term for a negative ion is an anion. Simple ions are formed when atoms lose or gain electrons. A molecular ion is an ion which contains more than one atom. Some examples of simple ions and molecular ions are given in the table below. Exam Tip The molecular ions which you need to know are listed on the back of the Data Leaflet which is supplied with your exam. Simple ions Molecular ions Name of ion Formula of ion Name of ion Formula of ion Sodium ion Na+ Ammonium ion NH4+ Calcium ion Ca2+ Hydroxide ion OH- Oxide ion O2- Nitrate ion NO3- Chloride ion Cl- Carbonate ion CO32- Nitride ion N3- Sulfate ion SO42- “Bananas contain potassium” is a comment often made in the media and by nutritionists but in truth bananas contain potassium ions. pg 16  GCSE CHEMISTRY: GUIDANCE Bonding Dot and cross diagrams A dot and cross diagram is used to show movement or sharing of electrons. For the formation of an ionic compound dots () are used to represent electrons in one atom and crosses (×) are used to represent electrons in another atom. This allows the movement of electrons to be easily observed. The formation of ionic compounds is limited to atoms of elements in Groups 1 (I) and 2 (II) with atoms of elements in Group 6 (VI) and 7 (VII). 1. Formation of sodium chloride Atoms of sodium react with atoms of chlorine to form sodium chloride. The electronic configuration of a sodium atom is 2, 8, 1. The electronic configuration of a chlorine atom is 2, 8, 7. When metal atoms react with non-metal atoms, electrons are often transferred from the metal atom to the non-metal atom so that the ions formed are more stable as they have a full outer shell of electrons. ×× ×× × × × ×× ×× sodium atom, Na chlorine atom, Cl ×× ×× × × × × × ×× sodium ion, Na+ chloride ion, Cl- The sodium atom loses its outer shell electron and becomes Na+. The chlorine atom gains one outer shell electron and becomes a chloride ion, Cl-. The ratio of sodium ions to chloride ions is 1:1 and therefore the formula of sodium chloride is NaCl (1Na and 1Cl). pg 17  GCSE CHEMISTRY: GUIDANCE Bonding 2. Formation of magnesium fluoride Atoms of magnesium react with atoms of fluorine to form magnesium fluoride. The electronic configuration of a magnesium atom is 2, 8, 2. The electronic configuration of a fluorine atom is 2, 7. ×× ×× ×× ×× × × × ×× × × × × × ×× ×× × magnesium atom, Mg fluorine atom, F magnesium ion, Mg2+ fluoride ion, F- × fluorine atom, F fluoride ion, F- The magnesium atom loses its two outer shell electrons and becomes Mg2+. The two fluorine atoms each gain one outer shell electron and become fluoride ions, F-. The ratio of magnesium ions to fluoride ions is 1:2 and therefore the formula of magnesium fluoride is MgF2 (1Mg and 2F). Exam Tip The symbol for fluorine is F and not Fl which is a common mistake. Always use the data leaflet as it contains a Periodic Table if you are not sure of a symbol for an element. pg 18  GCSE CHEMISTRY: GUIDANCE Bonding WORKED EXAMPLE Draw a dot and cross diagram to show how atoms of magnesium react with atoms of oxygen to form magnesium oxide. Answer: ×× ×× ×× ×× × × × ×× × × × × × ×× ×× × × magnesium atom, Mg oxygen atom, O magnesium ion, Mg2+ oxide ion, O2- Explanation: The electronic configuration of a magnesium atom is 2, 8, 2. The electronic configuration of an oxygen atom is 2, 6. A magnesium atom loses its two outer shell electrons to become a magnesium ion, Mg2+. An oxygen atom gains two electrons into its outer shell and becomes an oxide ion, O2-. The ratio of Mg2+: O2- is 1:1 so the formula of magnesium oxide is MgO. Svante Arrhenius was the first chemist to propose that an ion was an atom carrying a positive or negative charge. He was also the first to describe the Greenhouse effect. The climate change activist, Greta Thunberg, is a descendant of Arrhenius. Test yourself 1.2.1 1. Write electronic configurations to explain how calcium atoms react with chlorine atoms to form calcium chloride. Give the charges of the ions formed. 2. Explain what is meant by ionic bonding. 3. Draw a dot and cross diagram to show how atoms of lithium react with atoms of oxygen to form lithium oxide. pg 19  GCSE CHEMISTRY: GUIDANCE Bonding Covalent bonding A covalent bond is formed when electrons are shared between atoms. Covalent bonding is typical of non-metallic elements and compounds. Exam Tip The atoms of some non-metallic elements are so reactive that they cannot exist as atoms and so share electrons and form covalently bonded molecules of the element. All the atoms are still the same so they are still elements, but they exist as molecules. Covalent bonds are strong and require substantial energy to break them. Covalent bonds may be single, double or triple bonds. A single covalent bond between two atoms is a shared pair of electrons, a double covalent bond is two pairs of electrons shared between two atoms and a triple covalent bond is 3 pairs of electrons shared between two atoms. A dot and cross diagram can be used to show the outer shell electrons in the atoms and how the electrons are shared. A molecule is a particle that consists of two or more atoms chemically bonded together. A single covalent bond may be represented using a line ( ) between the symbols for the atoms. A double covalent bond is represented by two lines ( ) and a triple covalent bond as three lines between the symbols for the atoms ( ). The table below contains some examples of covalently bonded molecules containing single, double and triple covalent bonds and a representation of the bonding using lines. pg 20  GCSE CHEMISTRY: GUIDANCE Bonding Name Formula Bonding diagram Dot and cross diagram Hydrogen H2 H–H HxH xx x Chlorine Cl2 Cl–Cl x Cl x Cl xx Hydrogen chloride HCl H–Cl Hx Cl xx Water H2O O H x O xx –– x H H H xx Ammonia NH3 N Hx N xH –– x – H HH H H H x Hx – – Methane CH4 H–C–H C xH x H H x Oxygen O2 O O OxO x Nitrogen N2 N N x N xN x x x x Carbon dioxide CO2 O C O Ox C x O pg 21  GCSE CHEMISTRY: GUIDANCE Bonding The dots and crosses represent outer shell electrons. The × between atoms represents a shared pair of electrons often called a bonding pair of electrons. One electron in the bond comes from each atom. The pairs of outer electrons which are not involved in bonding are called lone pairs of electrons. xx lone pairs of electrons Hx O x x in a water molecule x H Elements like hydrogen, chlorine, nitrogen and oxygen exists as molecules. They are known as diatomic elements as they form molecules containing two atoms. In equations, these elements are written as the molecule such as H2, Cl2, N2 and O2. There are 7 diatomic elements you have to remember: H2, N2, O2 and the Group 7 (VII) elements, F2, Cl2, Br2, I2. O CH3 H3C C N N C C H C C N O N The bonding diagram shows the molecule caffeine. It contains covalent bonds represented by lines. A typical cup of coffee contains between 80-100 mg of CH3 caffeine. Metallic bonding The bonding in metals is described as metallic bonding. Metallic bonding is the attraction between positive ions in a regular lattice and delocalised electrons. The diagram below shows the positive ions and the delocalised electrons in the lattice for sodium metal. delocalized electron e- e- e- e- Na+ Na+ Na+ Na+ Na+ e- e- e- e- Na+ Na+ Na+ Na+ Na+ e- e- e- e- e- e- e- Na+ Na+ Na+ Na+ Na+ e- e- e- pg 22  GCSE CHEMISTRY: GUIDANCE Bonding A lattice is a regular arrangement of particles in three dimensions. Delocalised electrons are the outer shell electrons of atoms in the metal which are free to move through the structure. Test yourself 1.2.2 1. Name two diatomic elements. 2. Draw a dot and cross diagram to show the bonding in a molecule of ammonia, NH3. 3. How many lone pairs are in a molecule of chlorine, Cl2? 4. What is meant by metallic bonding? 5. A bonding diagram of a molecule is shown below. H H O H C C C OH H H (a) What is represented by the lines in the bonding diagram? (b) How many hydrogen atoms are present in the molecule? 6. A list of some molecules containing covalent bonds is shown below. ammonia carbon dioxide chlorine hydrogen hydrogen chloride methane nitrogen oxygen water (a) Which of the molecules contain double covalent bonds? (b) Which of the molecules contain no lone pairs of electrons? (c) Which of the molecules are compounds? (d) Which of the molecules contains four single covalent bonds? (e) Draw dot and cross diagrams for the following molecules. (i) hydrogen (ii) methane (iii) carbon dioxide (iv) nitrogen pg 23  GCSE CHEMISTRY: GUIDANCE 1.3 Structures This section will look at the different types of structures and related properties – ionic, molecular covalent, giant covalent and metallic as well as focussing on alloys and the allotropes of carbon including graphene. Mathematical content Students will calculate the percentage of gold based on carat ratings as well as work with melting and boiling point data in °C. Learning outcomes Ionic structures use the accepted structural model for giant ionic lattices to explain the physical properties of ionic substances such as sodium chloride, including melting point, boiling point and electrical conductivity (drawing a diagram of a giant ionic lattice is not expected but students should be able to recognise it) recall that most ionic compounds are soluble in water Molecular covalent structures use the accepted structural model for molecular covalent structures to explain the physical properties of molecular covalent structures such as iodine and carbon dioxide, including melting point, boiling point and electrical conductivity demonstrate knowledge and understanding that the intermolecular forces between covalent molecules are weak forces called van der Waals’ forces recall that many covalent molecular substances are insoluble in water Giant covalent structures demonstrate knowledge and understanding of the giant covalent structure of carbon (diamond) and carbon (graphite), and predict and explain their physical properties, including: – electrical conductivity – hardness – melting point and boiling point – their uses in cutting tools (diamond), lubricants and pencils (graphite) pg 24  GCSE CHEMISTRY: GUIDANCE Structures Metallic structures use the accepted structural model for metals to predict and explain their structure and physical properties including melting point, malleability, ductility and electrical conductivity demonstrate knowledge and understanding that an alloy is a mixture of two or more elements, at least one of which is a metal, and the resulting mixture has metallic properties demonstrate knowledge and understanding that the different sizes of atoms in an alloy distort the layers in the metallic structure, making it more difficult for them to slide over each other, and so alloys are harder than pure metals recall that gold used in jewellery is usually an alloy with silver, copper and zinc, that the proportion of gold is measured in carats, and that 24 carat gold indicates pure gold and 18 carat gold indicates 75% gold Structure and bonding of carbon demonstrate knowledge and understanding that carbon can form four covalent bonds demonstrate knowledge and understanding of the structure of graphene (a single atom thick layer of graphite), explain its physical properties, including strength and electrical conductivity, and recall its uses such as those in batteries and solar cells demonstrate knowledge and understanding of the meaning of the term allotrope as applied to carbon (diamond), carbon (graphite) and graphene Classification of structures use given information to classify the structure of substances as giant ionic lattice, molecular covalent, giant covalent or metallic pg 25  GCSE CHEMISTRY: GUIDANCE Structures Ionic structures Ionic compounds such as sodium chloride and magnesium oxide exist as an ionic lattice when they are solids. An ionic lattice consists of alternate positive and negative ions in a regular arrangement. The diagram below shows the ionic lattice for sodium chloride. Key positive Na+ ion negative Cl- ion Exam Tip You do not have to be able to draw the lattice just understand that it is alternate positive and negative ions in all directions. The attractions between the oppositely charged ions is ionic bonding. Ionic bonding is strong and requires substantial energy to break it. Also, when solid, the ions are held in the lattice and cannot move. The properties of ionic compounds depend on the bonding (ionic bonding) and the structure (ionic lattice). The table below shows some of the properties of ionic compounds with an explanation based on the bonding and/or structure. Property Explanation 1. Ionic compounds have high melting points The ionic bonding is strong and and boiling points/they are solids at room substantial energy is required to temperature break it 2. Ionic compounds cannot conduct electricity The ions cannot move or carry when solid charge 3. Ionic compounds can conduct electricity when The ions can move and carry molten or when dissolved in water charge You do not have to explain this 4. Most ionic compounds are soluble in water property pg 26  GCSE CHEMISTRY: GUIDANCE Structures Test yourself 1.3.1 1. What is an ionic lattice? 2. Explain why ionic compounds conduct electricity when dissolved in water. 3. Name two ionic compounds. 4. Explain why ionic compounds have high melting points. Molecular covalent structures Most non-metallic compounds and non-metallic elements exist as simple molecules. (Exceptions to this are carbon and silicon which exist as giant covalent structures and the Group 0 elements which exist as atoms.) The molecules contain covalent bonding and their structure is described as molecular covalent as they exist as simple molecules. There are weak forces of attraction between the molecules which are called van der Waals’ forces. Covalent bonding is strong and requires substantial energy to break but van der Waals’ forces between the molecules are weak and little energy is required to break them. Some examples of molecular covalent substances are the molecules listed in the table in Section 1.2 Covalent bonding. These are H2, Cl2, HCl, H2O, NH3, CH4, O2, CO2, N2. All the diatomic elements are molecular covalent substances. The table below shows some of the properties of molecular covalent substances with an explanation based on the bonding and/or structure. Property Explanation 1. Molecular covalent substances have low The van der Waals’ forces between melting points and boiling points/they are the molecules are weak and gases, liquids or low melting point solids at require little energy to break room temperature 2. Molecular covalent substances cannot conduct They do not contain charged electricity particles 3. Most molecular covalent substances are You do not have to explain this insoluble in water property When water in a kettle boils, the van der Waals’ forces between the water molecules are broken but the covalent bonds in the molecules are not broken. The water vapour formed still contains water molecules. pg 27  GCSE CHEMISTRY: GUIDANCE Structures Giant covalent structures There are three giant covalent structures which we examine: carbon (diamond), carbon (graphite) and carbon (graphene). Carbon atoms can form a maximum of four bonds. Different forms of the same element in the same physical state are known as allotropes. Diamond, graphite and graphene are allotropes of carbon. A giant covalent structure is a regular arrangement of atoms held together by covalent bonds. The diagrams below show the structure and bonding in diamond and graphite. carbon (diamond) covalent bonds carbon atom carbon (graphite) covalent bond carbon atom weak bonds between layers carbon (graphene) covalent bond carbon atom pg 28  GCSE CHEMISTRY: GUIDANCE Structures The following gives details of the arrangement of the carbon atoms: Diamond: Each carbon atom is covalently bonded to 4 other carbon atoms The arrangement is described as tetrahedral. Graphite: Each carbon atom is covalently bonded to 3 other carbon atoms The arrangement is layers of hexagons with weak bonds between the layers caused by delocalised electrons. Graphene: A single layer of graphite, one atom thick. The properties of diamond, graphite and graphene depend on the bonding and structure. Property Explanation Diamond 1. Does not contain free ions or 1. Does not conduct electricity electrons to move and carry 2. Hard charge 3. High melting point and boiling point 2. Strong covalent bonds and a rigid 3-dimensional structure 3. Many strong covalent bonds which require substantial energy to break Graphite 1. Contains delocalised electrons 1. Conducts electricity which can move and carry 2. Soft charge 3. High melting point and boiling point 2. Weak bonds between the layers so the layers can slide off 3. Many strong covalent bonds which require substantial energy to break Graphene 1. Contains delocalised electrons 1. Conducts electricity which can move and carry 2. Strong charge 3. High melting point and boiling point 2. Contains many strong covalent 4. Transparent bonds 3. Contains many strong covalent bonds which require substantial energy to break 4. Single atom thick layer of carbon atoms pg 29  GCSE CHEMISTRY: GUIDANCE Structures Uses of allotropes of carbon Diamond is used in cutting tools such as drill heads for drilling through rock as it is the hardest naturally known substance. Graphite is used as a lubricant and in pencils as the layers can slide over each other and off onto a piece of paper when used in a pencil. Graphene is used in batteries as it can conduct electricity and is very light and also in solar cells again because it can conduct electricity but also because it is transparent. Test yourself 1.3.2 1. Explain why chlorine is a gas at room temperature. 2. From the list below choose two molecular covalent substances: hydrogen chloride calcium chloride sodium chloride aluminium oxide copper(II) oxide carbon dioxide 3. What is meant by the term allotropes? 4. Name two allotropes of carbon. 5. Describe the structure, bonding and arrangement of atoms in diamond. 6. State one use of graphite. 7. Explain why graphene is used in solar cells. pg 30  GCSE CHEMISTRY: GUIDANCE Structures Metallic structures Metals contain metallic bonding which is the attraction between the positive ions and the delocalised electrons. The structure of a solid metal as a solid is known as a metallic lattice and it is a regular arrangement of positive ions held together by the “sea” of delocalised electrons. The properties of metals depend on the bonding (metallic bonding) and the structure (metallic lattice). The table below shows some of the properties of metals with an explanation based on the bonding and/or structure. Property Explanation The metallic bonding is strong 1. Metals have high melting points/they are and substantial energy is required solids at room temperature to break it The delocalised electrons can 2. Metals conduct electricity move and carry the charge The layers of positive ions can 3. Metals are malleable (they can be hammered slide over each other without into shape without breaking) disrupting the bonding The layers of positive ions can 4. Metals are ductile (they can be drawn out into slide over each other without wires) disrupting the bonding Alloys An alloy is a mixture of elements, at least one of which is metal, and the resulting mixture has metallic properties. Many pure metals are soft. The different sizes of atoms in the alloy distort the layers making it harder for the layers to slide over each other. This makes alloys harder than pure metals. However the alloys still have the other properties of metals for example electrical conductivity and are often lighter (duralumin used for aircraft bodies) and more resistant to corrosion such as stainless steel. Gold used in jewellery is often an alloy of gold with other metals. 24 carat gold is pure gold. 18 carat gold is 75% pure gold. The percentage is worked out using 24 carat gold as 100% pure gold. carat rating % pure gold = × 100 24 pg 31  GCSE CHEMISTRY: GUIDANCE Structures Worked example Calculate the carat rating of jewellery gold which contains 37.5 % pure gold. carat rating 37.5 = × 100 24 0.375 × 24 = carat rating = 9 carats Stainless steel is an alloy of iron containing chromium and carbon. The chromium and carbon in the alloy prevent the iron from rusting. Classification of structures If you are given information on different substances, it is possible to classify them as one of the four types of structures: ionic, molecular covalent, giant covalent or metallic. The table below summarises the general properties. Melting point/ Electrical Solubility in Other Structure Boiling point conductivity water properties Ionic High Do not conduct as Generally lattice solid but conduct soluble when molten or dissolved in water Molecular Low Do not conduct Generally covalent insoluble Giant Very high Diamond: does not Insoluble Diamond: Hard covalent conduct Graphite: Soft Graphite and Graphene: graphene: conduct Strong Metallic High Conducts Generally Malleable and lattice insoluble but ductile some metals react with water Specific data may be given and you may have to classify substances based on their properties. pg 32  GCSE CHEMISTRY: GUIDANCE Structures Test yourself 1.3.3 1. Explain why iron has a high melting point. 2. What is meant by the terms below? (a) malleable (b) ductile (c) alloy 3. A substance melts at 714 °C and does not conduct electricity when solid. It is soluble in water and the solution formed does conduct electricity. Choose a substance from the list which could be this substance and explain your answers in terms of the structure and bonding. aluminium carbon dioxide copper diamond magnesium chloride water 4. Calculate the percentage of pure gold in an alloy labelled as 15 carat. 5. State the bonding and structure found in each of the following. (a) sodium (b) graphite (c) hydrogen (d) iron(II) oxide pg 33  GCSE CHEMISTRY: GUIDANCE 1.4 Nanoparticles This unit looks at the size, properties and uses of nanoparticles compared to traditional bulk materials. Mathematical content Students will work with units of measurement in the 10-9 range. 1 nm = 1 × 10-9 m. Calculations on surface area to volume ratios of cubes of different side lengths are carried out. Learning outcomes demonstrate knowledge and understanding that nanoparticles are structures that are 1-100 nm in size and contain a few hundred atoms demonstrate knowledge and understanding of surface area to volume relationships and that, as the side of a cube decreases by a factor of 10, the surface area to volume ratio increases by a factor of 10 demonstrate knowledge and understanding that nanoparticles have properties different from those for the same material in bulk, due to their high surface area to volume ratio evaluate the benefits of nanoparticles in sun creams, including better skin coverage and more effective protection from the Sun’s ultraviolet rays, and the risks, such as potential cell damage in the body and harmful effects on the environment Nanometres Nanoparticles are particles which are in the range of 1 – 100 nm in size and contain a few hundred atoms. 1 nm is 1 nanometre and is 1 × 10-9 m. This means there are 1 million nanometres in 1 millimetre. Surface area to volume ratio The surface area to volume ratio of a cube is the simplest ratio of the total surface area of all 6 sides of the cube to the overall volume. Two cubes with side lengths 20 mm and 2 mm are shown below. 20 mm 2 mm pg 34  GCSE CHEMISTRY: GUIDANCE Nanoparticles For the cube with side length 20 mm. Total area = 6 × 20 × 20 = 2400 mm2 Volume = 20 × 20 × 20 = 8000 mm3 3 Surface area to volume ratio = 2400 : 8000 = 3:10 = = 0.3 10 For the cube with side length 2 mm Total area = 6 × 2 × 2 = 24 mm2 Volume = 2 × 2 × 2 = 8 mm3 3 Surface area to volume ratio = 24 : 8 = 3:1 = = 3 1 As the length of the side of a cube decreases by a factor of 10, the surface area to volume ratio increases by a factor of 10. This is important as smaller particles have a much greater surface area for the same volume and this gives them different properties to bulk materials which have larger particles. Uses of nanoparticle materials Nanoparticles are used in sun creams as they give better skin coverage, rub on clear and give better protection from the ultraviolet rays from the Sun. However, there are also risks associated with the use of nanoparticles as their effects on the human body are unknown. They may cause cell damage and as they wash off from our skin, they may be harmful for the environment. Suncream advertised as without nanoparticles (“sans nanoparticles”) due to concerns about their effects. Test yourself 1.4.1 1. What is a nanometre? 2. What is the size range of a nanoparticle? 3. A cubic particle has a side length of 15 mm. A second particle has a side length of 15 cm. Calculate the change in the surface area to volume ratio. 4. State one reason why nanoparticles are used in sun creams. pg 35 1.5 Symbols, formulae and equations This unit looks at writing chemical formula for elements and compounds as well as the different types of equations which can be written to represent chemical reactions. Mathematical content Students will use charges to work out the formula of ionic compounds and balancing numbers to balance the atoms in equations for chemical reactions. Learning outcomes recognise symbols and names for common elements and recall the diatomic elements interpret chemical formulae by naming the elements and stating the number of each type of atom present write chemical formulae of compounds demonstrate understanding that chemical reactions use up reactants and produce new substances called products construct word equations to describe the range of reactions covered in this specification recognise that in a chemical reaction no atoms are lost or made but they are rearranged, and as a result we can write balanced symbol equations showing the atoms involved write balanced symbol equations for all reactions covered in this specification and for unfamiliar chemical reactions when the names of the reactants and products are specified write balanced ionic equations for reactions covered in this specification write half equations for reactions covered in this specification demonstrate knowledge and understanding that in chemical equations the three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions, and include appropriate state symbols in equations for the reactions in this specification pg 36 Symbols, formulae and equations Symbols and formulae of elements For metallic elements and most non-metallic elements including boron, carbon, silicon, phosphorus, sulfur and the Group 0 elements, the symbol for the element is used to represent the element in chemical equations. Exam Tip This is not strictly true as at A-level you will learn more about phosphorus and sulfur as they exist as P4 and S8 molecules but for equations at this level, P and S are acceptable. For the diatomic elements, the symbol with a subscript 2 is used to represent these elements in equations, for example H2, N2, O2, F2, Cl2, Br2 and I2. Exam Tip There are 7 diatomic elements to remember. Make a mnemonic to remember them. Most compounds you will come across during the course are ionic compounds and you can use the charges on the ions involved to determine the formula. Other covalent compounds must be learned but you will use these formulae a lot as you work through the course. Group in Periodic Table Charge on ion Examples Group 1 (I) + Li+, Na+, K+ Group 2 (II) 2+ Mg2+, Ca2+ Group 3 (III) 3+ Al3+ Group 4 (IV) Do not generally form ions None Group 5 (V) 3- N3-, P3- Group 6 (VI) 2- O2-, S2- Group 7 (VII) - F-, Cl-, Br-, I- Group 0 (VII) Unreactive None pg 37 Symbols, formulae and equations The transition metals are the block of elements between Group 2 and Group 3 and these elements form ions with different charges. In general, for a transition metal, the charge is given after the name of the metal. For example, iron(III) oxide would contain the Fe3+ ion; copper(II) chloride would contain the Cu2+ ion; silver(I) nitrate would contain the Ag+ ion. If no number appears in the name of a compound of one of these elements, you can assume it is the 2+ ion. Apart from silver which is usually Ag+ and should be written as silver(I) but often this is left out and silver is just written. Note that some of the metals from Groups 3 to 6 can also form ions with different charges and the charge may be given, for example lead(II) is Pb2+. Exam Tip Some of these ions are on the back of your Data Leaflet but you should be able to work with unusual ones as well such as copper(I) oxide where Cu+ and O2- give Cu2O. Molecular ions A molecular ion is a charged particle containing more than one atom. The molecular ions you need to know are given on the back of your Data Leaflet together with some transition metal ions. The information from the back of the Data Leaflet is shown below. Positive ions Negative ions Name Symbol Name Symbol Ammonium NH4+ Butanoate C3H7COO- Chromium(III) Cr3+ Carbonate CO2- 3 Copper(II) Cu2+ Dichromate Cr2O2- 7 Iron(II) Fe2+ Ethanoate CH3COO- Iron(III) Fe3+ Hydrogencarbonate HCO3- Lead(II) Pb2+ Hydroxide OH- Silver Ag+ Methanoate HCOO- Zinc Zn2+ Nitrate NO3- Propanoate C2H5COO- Sulfate SO42- Sulfite SO32- pg 38 Symbols, formulae and equations Formulae of compounds Some rules: 1. Numbers of ions in a compound are shown using a subscript number after the ion. No number is needed where 1 of an ion is required. 2. The overall charge of a compound should be zero. 3. Where two or more molecular ions are used they need a bracket with a subscript number outside the right of the bracket. 4. Covalent compounds need to be learned. Examples: 1. What is the formula of sodium chloride? sodium is in Group 1 so its ion is Na+ chlorine is in Group 7 so its ion is Cl- To make a compound with no overall charge we need one Na+ and one Cl- so the formula of sodium chloride is NaCl. 2. What is the formula of magnesium oxide? Mg Group 2 so Mg2+; O Group 6 so O2- 1 Mg2+ and 1 O2- so MgO (total 2+ and 2-) 3. What is the formula of calcium fluoride? Ca Group 2 so Ca2+; F Group 7 so F- 1 Ca2+ and 2 F- so CaF2 (total 2+ and 2-) 4. What is the formula of aluminium oxide? Al Group 3 so Al3+; O Group 6 so O2- 2Al3+ and 3O2- so Al2O3 (total 6+ and 6-) 5. What is the formula of lithium hydroxide First example with a molecular ion, hydroxide is OH- Li Group 1 so Li+; hydroxide is OH- 1 Li+ and 1 OH- so LiOH (1+ and 1- and no bracket needed for OH-) 6. What is the formula of copper(II) hydroxide? Cu transition metal but (II) so Cu2+; hydroxide is OH- 1Cu2+ and 2OH- so Cu(OH)2 (2+ and 2- and bracket needed as 2 × OH-) 7. What is the formula of silver nitrate? Ag transition metal and normally silver(I) so Ag+; nitrate is NO3- 1Ag+ and 1NO3- so AgNO3 (1+ and 1- and no bracket required) pg 39 Symbols, formulae and equations 8. What is the formula of chromium(III) sulfate? Cr transition metal and (III) given so Cr3+; sulfate is SO42- 2 Cr3+ and 3SO42- so Cr2(SO4)3 (6+ and 6- and brackets required for SO2- 4) 9. What is the formula of ammonium carbonate? ammonium is NH4+; carbonate is CO32- 2NH4+ and 1CO32- so (NH4)2CO3 (2+ and 2- and brackets required for NH4+) 10. What is the formula of copper(II) sulfate? Cu transition metal and (II) given so Cu2+; sulfate is SO42- 1 Cu2+ and 1 SO42- so CuSO4 (2+ and 2-) Covalent compounds The table below shows some covalent compounds. Compounds Formula Compound Formula Ammonia NH3 Hydrogen chloride HCl Water H2O Hydrochloric acid HCl Methane CH4 Nitric acid HNO3 Carbon dioxide CO2 Sulfuric acid H2SO4 Carbon monoxide CO Ethanol C2H5OH Sulfur dioxide SO2 Ethanoic acid CH3COOH Hydrogen peroxide H2O2 Sulfurous acid H2SO3 Exam Tip The formula of hydrogen chloride and hydrochloric acid are the same as hydrochloric acid is a solution of hydrogen chloride in water. Hydrogen chloride is one of the molecular covalent substances which is soluble in water. More covalent compounds will be added to the ones above especially in the organic section 2.5. pg 40 Symbols, formulae and equations Number of atoms in a formula It is important to be able to count the number of atoms in the formula for a compound. The formula of lithium hydroxide is LiOH and it contains 1 Li atom, 1 O atom and 1 H atom. The formula of iron(III) nitrate is Fe(NO3)3 and it contains 1 Fe atom, 3 N atoms and 9 O atoms. The formula of potassium dichromate is K2Cr2O7 and it contains 2 K atoms, 2 Cr atoms and 7 O atoms. Exam Tip Note the number of atoms inside the brackets is multiplied by the number outside the bracket. Test yourself 1.5.1 1. Write symbols (or formulae for diatomic elements) for the following. (a) iron (b) chlorine (c) magnesium (d) nitrogen (e) argon 2. Write formulae for the following compounds. (a) potassium oxide (b) copper(II) chloride (c) water (d) chromium(III) nitrate (e) manganese(IV) oxide (f) ammonia 3. Name the following compounds. (a) LiOH (b) CaSO4 (c) CH4 (d) Fe(NO3)2 (e) Cr2O3 (f) KHCO3 (g) Al2(SO4)3 (h) (NH4)2SO4 Balanced symbol equations In a chemical reaction, reactants are changed into products. No atoms are lost or made during a chemical reaction. The total number of atoms in the reactants must equal the total number of atoms in the products. A balanced symbol equation shows the formulae of the reactants on the left with an arrow and the formulae of the products on the right. Some rules: 1. Never change a formula to balance an equation. 2 Numbers can only be added before a formula. 3. A number added before a formula multiplies all the atoms in the formula by that number. 4. 3 marks are awarded to an equation which requires balancing numbers. 5. 2 marks are awarded to an equation which does not require balancing numbers. 6. Beware of state symbols as they are an extra mark and they change a 2 mark equation (which does not require balancing numbers) into a 3 mark question. pg 41 Symbols, formulae and equations Exam Tip Watch out for the number of marks for a balanced symbol equation as this is an indication of what is required. Examples: Write balanced symbol equations for the reactions shown by the word equations below. 1. sodium + chlorine → sodium chloride Answer: Write the formula for the reactants and products in the style of an equation using an arrow: Na + Cl2 → NaCl Note that chlorine is Cl2; Na+ and Cl- gives NaCl for sodium chloride. The unbalanced equation is worth 2 marks. On the left there is 1 Na and 2 Cl in Cl2; on the right there is 1 Na and 1 Cl in NaCl. Put a 2 in front of NaCl. Na + Cl2 → 2NaCl This gives 2 Cl on the right but there are now also 2 Na on the left. Put a 2 in front of Na on the left. 2Na + Cl2 → 2NaCl This is the balanced equation and is worth 3 marks. 2. calcium carbonate → calcium oxide + carbon dioxide Answer: This is worth 2 marks so there are no balancing numbers required. CaCO3 → CaO + CO2 Note that calcium carbonate is formed from Ca2+ and CO2- 3 so is CaCO3; calcium oxide 2+ 2- is formed from Ca and O so is CaO; carbon dioxide is a covalent compound from the table. 1 Ca on left and on right; 1 C on left and on right; 3 O on left and on right so this is the balanced equation and is worth 2 marks. Exam Tip Note CaCO3 means 1Ca, 1C and 3O. The small number 3 only relates to the element immediately before it unless it comes after a bracket. 3. calcium + water → calcium hydroxide + hydrogen Answer: Unbalanced equation: Ca + H2O → Ca(OH)2 + H2 The bracket around OH means that there are 2 O and 2 H in Ca(OH)2 as well as the 1 Ca Balanced equation: Ca + 2H2O → Ca(OH)2 + H2 pg 42 Symbols, formulae and equations 4. aluminium + oxygen → aluminium oxide Answer: Unbalanced equation: Al + O2 → Al2O3 Balanced equation: 4Al + 3O2 → 2Al2O3 5. lead(II) nitrate + potassium iodide → lead(II) iodide + potassium nitrate Answer: Unbalanced equation: Pb(NO3)2 + KI → PbI2 + KNO3 Balanced equation: Pb(NO3)2 + 2KI → PbI2 + 2KNO3 The bracket around the NO3 means 2 NO3 (or 2 N and 6 O). As the NO3 is not changing it is easier to count how many NO3 there are on both the left and the right. However this is not always the case as if the NO3 were to break up the individual atoms must be considered. Exam Tip You will practise writing balanced symbol equations throughout the rest of the course so make sure you can write them and balance them. State symbols State symbols are used to show the physical state of the reactants and products in a chemical reaction. The state symbols used are: (s) represents solid (l) represents liquid (g) represents gas (aq) represents aqueous (i.e. dissolved in water) Dilute acids and dissolved ionic compounds use the (aq) state symbol. The explosive in dynamite has the chemical formula C7H5N3O6 and when the solid explodes it releases a lot of heat and forms soot (carbon) and a lot of gas which expands quickly causing the force of the blast. The equation for the explosion with state symbols is: 2C7H5N3O6(s) → 7CO(g) + 7C(s) + 5H2O(g) + 3N2(g) pg 43 Symbols, formulae and equations Ionic equations and half equations An ionic equation shows how ions react together leaving out any ions which do not take part in the reaction. Ionic equations are often used for neutralisation reactions and also for displacement and precipitation reactions. The ionic equation for any acid reacting with an alkali is H+ + OH- → H2O. The H+ ions from the acid react with OH- from the alkali to product water, H2O. Any other ions in the solution from the acid or alkali remain in the solution. Exam Tip This will be important in section 1.8 in acid reactions and in salt preparation. For the following displacement reaction: magnesium + copper(II) sulfate → magnesium sulfate + copper Magnesium displaces copper. However, looking at all the atoms and ions involved shows that one ion does not take part in the reaction. Reactants Products Mg + CuSO4 → MgSO4 + Cu Cu2+ ions Mg2+ ions Mg atoms SO42- ions SO42- ions Cu atoms The sulfate ions, SO2- 4 , do not take part in the reaction as they are unchanged and so they can be left out of the ionic equation. The ionic equation is: Mg + Cu2+ → Mg2+ + Cu Half equations are equations involving electrons and they can be used to show how the electrons are moving during the reaction. There are two half equations which can be written based on this ionic equation: Mg → Mg2+ + 2e- Cu2+ + 2e- → Cu Each magnesium atom is losing two electrons to become Mg2+ and this is shown as + 2e- on the right hand side. Each copper(II) ion is gaining two electrons to become Cu and this is shown as +2e- on the left hand side. Exam Tip Remember that ionic equations do not show electrons and half equations do. Two half equations can form an ionic equation if the electrons cancel out. pg 44 Symbols, formulae and equations Test yourself 1.5.2 1. Write balanced symbol equations for the following reactions. Note these are all mark equations except (d) which is worth. (a) magnesium + oxygen → magnesium oxide (b) lithium + oxygen → lithium oxide (c) aluminium + bromine → aluminium bromide (d) carbon + oxygen → carbon dioxide 2. Balance the following symbol equations. (a) Na + H2O → NaOH + H2 (b) Mg + HCl → MgCl2 + H2 (c) Al + ZnSO4 → Al2(SO3)3 + Zn (d) FeCl3 + NaOH → Fe(OH)3 + NaCl 3. Write balanced symbol equation for the following reactions. Note (b) is worth marks but the rest are (a) calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide (b) calcium hydroxide + carbon dioxide → calcium carbonate + water (c) aluminium oxide + nitric acid → aluminium nitrate + water (d) chlorine + potassium iodide → potassium chloride + iodine 4. Add state symbols to the following reactions. (a) Solid magnesium reacts with dilute sulfuric acid forming a solution of magnesium sulfate and releasing hydrogen gas. Mg( ) + H2SO4( ) → MgSO4( ) + H2( ) (b) Solid potassium reacts with water forming potassium hydroxide in solution and hydrogen gas. 2K( ) + 2H2O( ) → 2KOH( ) + H2( ) (c) Ammonia gas and hydrogen chloride gas form solid ammonium chloride. NH3( ) + HCl( ) → NH4Cl( ) 5. Write ionic equation for the following reactions: (a) Zn + CuSO4 → ZnSO4 + Cu (b) NaOH + HCl → NaCl + H2O (c) 2AgNO3 + Cu → Cu(NO3)2 + 2Ag 6. Write the half equations for the two processes in 6(c). pg 45  GCSE CHEMISTRY: GUIDANCE 1.6 The Periodic Table This section looks at the Periodic Table and its history and organisation including details of the physical and chemical properties of Group 1(I), the transition metals, Group 7(VII) and Group 0. Learning outcomes Basic structure of the Periodic Table describe how Mendeleev arranged the elements in the Periodic Table and left gaps for elements that had not been discovered at that time, and how this enabled him to predict properties of undiscovered elements demonstrate knowledge and understanding of how scientific ideas have changed over time in terms of the differences and similarities between Mendeleev’s Periodic Table and the modern Periodic Table describe an element as a substance that consists of only one type of atom and demonstrate understanding that elements cannot be broken down into simpler substances by chemical means demonstrate knowledge and understanding that a group is a vertical column in the Periodic Table and a period is a horizontal row identify and recall the position of metals and non-metals in the Periodic Table and distinguish between them according to their properties, including conduction of heat and electricity, ductility, malleability, melting point and sonority identify elements as solids, liquids and gases (at room temperature and pressure) in the Periodic Table demonstrate knowledge and understanding that elements in the same group in the Periodic Table have the same number of electrons in their outer shell and this gives them similar chemical properties recall that elements with similar properties appear in the same group (for example Group 1 (I) and Group 2 (II) are groups of reactive metals, Group 7 (VII) is a group of reactive non-metals and Group 0 is a group of non-reactive non-metals), locate these groups in the Periodic Table and recall the names of the groups pg 46  GCSE CHEMISTRY: GUIDANCE The Periodic Table Group 1 (I) demonstrate knowledge and understanding that the alkali metals have low density and the first three are less dense than water. assess and manage risks associated with the storage and use of alkali metals and recall that alkali metals are easily cut, are shiny when freshly cut and tarnish rapidly in air. demonstrate knowledge and understanding that Group 1 (I) metals react with water to produce hydrogen and a metal hydroxide, and give observations for the reactions. demonstrate knowledge and understanding that alkali metals have similar chemical properties because when they react an atom loses an electron to form a positive ion with a stable electronic configuration. write half equations for the formation of a Group 1 (I) ion from its atom. demonstrate knowledge and understanding of how the trend in reactivity down the group depends on the outer shell electrons of the atoms. demonstrate knowledge and understanding that most Group 1 (I) compounds are white and dissolve in water to give colourless solutions. Group 7 (VII) recall data about the colour, physical state at room temperature and pressure, diatomicity and toxicity of the elements in Group 7 (VII), interpret given data to establish trends within the group and make predictions based on these trends. recall the observations when solid iodine sublimes on heating and demonstrate understanding of the term sublimation. describe how to test for chlorine gas (damp universal indicator paper changes to red and then bleaches white). investigate the displacement reactions of Group 7 (VII) elements with solutions of other halides to establish the trend in reactivity within the group and make predictions based on this trend. demonstrate knowledge and understanding of how the reactivity down the group depends on the outer shell electrons of the atoms. demonstrate knowledge and understanding that the halogens have similar chemical properties because when they react an atom gains an electron to form a negative ion with a stable electronic configuration. write half equations for the formation of a halide ion from a halogen molecule or atom. pg 47  GCSE CHEMISTRY: GUIDANCE The Periodic Table Group 0 use the concept of electronic configuration to explain the lack of reactivity and the stability of the noble gases. recall that the noble gases are colourless gases. demonstrate knowledge and understanding of the trend in boiling points of the noble gases going down the group. Transition metals compare the physical properties of the transition metals with Group 1 (I) elements, including melting point and density, and demonstrate understanding that the transition metals are much less reactive with water. demonstrate knowledge that transition elements form ions with different charges (for example iron(II) and iron(III)) and form coloured compounds: – copper(II) oxide is black; – copper(II) carbonate is green; – hydrated copper(II) sulfate is blue; and – copper(II) salts are usually blue in solution. Basic structure of the Periodic Table The Periodic Table was originally devised by Dmitri Mendeleev in 1867. There are several differences between the table devised by Mendeleev and the modern Periodic Table we use today. Some of the important differences are: Mendeleev’s table was arranged in order of atomic weight whereas the modern Periodic Table is in order of atomic number. Mendeleev’s table had gaps for undiscovered elements whereas the modern Periodic Table has no gaps. There were no noble gases in Mendeleev’s table but there are in the modern Periodic Table. There was no block of transition metals in Mendeleev’s table whereas there is in the modern Periodic Table. The gaps in Mendeleev’s Periodic Table of elements allowed him to predict the properties of the undiscovered elements using the trends in the Periodic Table. pg 48  GCSE CHEMISTRY: GUIDANCE The Periodic Table Dmitri Mendeleev and his Periodic System of Elements. Mendeleev’s Periodic Table: https://www.youtube.com/watch?v=fPnwBITSmgU The Periodic Table as supplied to you in your examinations in the Data Leaflet is shown below with some colour added for different groups. pg 49  GCSE CHEMISTRY: GUIDANCE The Periodic Table There are several important features to note from this Periodic Table: The table contains all known elements. An element is a substance which consists of only one type of atom and cannot be broken down into anything simpler by chemical means. A vertical column in the Periodic Table is called a group. There are 8 main groups which are labelled 1 – 7 and 0 in the Periodic Table shown. A horizontal row in the Periodic Table is called a period. Period 1 only contains two elements (hydrogen and helium) and is often missed when counting periods. The thick staircase line which starts below Boron separates the metals (to the left of the line) from the non-metals to the right of the line. Metals have characteristic metallic properties such as: – metals conduct heat and electricity – metals are ductile (can be drawn out into wires) – metals are malleable (can be hammered into shaped without breaking) – metals are sonorous (make a ringing sound when struck) Non-metals do not have these properties (except graphite and graphene which conduct electricity). Almost all the elements are solids at room temperature and pressure. There are 11 gases and 2 liquids at room temperature. The gaseous elements are: – hydrogen – nitrogen – oxygen – fluorine – chlorine – all the noble gases (helium, neon, argon, krypton, xenon and radon) The liquid elements are bromine (a non-metal) and mercury (a metal). Elements in the same group of the Periodic Table have the same number of electrons in their outer shell. Atoms of Group 1 elements have 1 electron in their outer shell. Atoms of Group 7 elements have 7 electrons in their outer shell. pg 50  GCSE CHEMISTRY: GUIDANCE The Periodic Table The main groups in the Periodic Table on page 49 are colour coded. – Group 1 is coloured pink on the left. It is a group of reactive metals. This group is called the alkali metals. – Group 2 is coloured orange on the left. It is also a group of reactive metals (but they are not as reactive as the metals in Group 1). This group is called the alkaline earth metals. – Group 7 is coloured blue on the right. It is a group of reactive non-metals. This group is called the halogens. – Group 0 is coloured light green on the right. It is a group of non-reactive non- metals. This group is called the noble gases. Hydrogen is coloured dark green at the top of the Periodic Table. Hydrogen does not have a true position in the Periodic Table as it shows a variety of properties. It is most often shown above the middle block. The middle block between Group 2 and 3 is known as the transition metals. Exam Tip The group numbers are most often written as standard numerals, 1, 2 etc but they are also written as roman numbers so be aware of these in other Periodic Tables I, II, III, IV, V, VI, VII and VIII. Group 0 can be called Group VIII. pg 51  GCSE CHEMISTRY: GUIDANCE The Periodic Table Test yourself 1.6.1 1. What name is given to the following groups of the Periodic Table? (a) Group 1 (b) Group 2 (c) Group 7 (d) Group 0 2. What elements are found in the following positions in the Periodic Table? (a) Period 3 and Group 2 (b) Period 1 and Group 0 (c) Period 4 and Group 7 (d) Period 2 and Group 6 3. Complete the table below. The first one has been done for you. Solid, liquid or gas at room Element Metal or non-metal temperature and pressure Name Symbol Metal Non-metal Solid Liquid Gas Sodium Na ✓ ✓ Hg Cl Iron Bromine Neon H pg 52  GCSE CHEMISTRY: GUIDANCE The Periodic Table Group 1 The Group 1 elements we consider are lithium, sodium and potassium. Group 1 metals are soft metals with low density. The first three (Li, Na and K) are less dense than water (and so float on water). Group 1 metals are highly reactive and are stored under oil. The metals are dull in appearance but show a shiny surface when freshly cut. This shiny surface changes to dull rapidly again as the metals react with oxygen and water vapour in the air. The shiny surface changing to dull is called tarnishing. The reactions of the Group 1 metals with oxygen form the metal oxide which are white solids. An example equation is: 4K + O2 → 2K2O The metals react vigorously with water. The reactivity increases down the group. When reacting a Group 1 metal with water, the metal is removed from the oil in which it is stored using tongs and a small piece is cut, the oil is removed and the metal is added to a large volume of water behind a safety screen. The general observations for lithium, sodium and potassium reacting with water are: – metal floats on the surface on the water – metal moves around the surface – fizzes – heat is released – a colourless solution is formed – the metal disappears For sodium and potassium, the metal also melts on the surface of the water forming a silvery ball. For potassium, there is a lilac flame and a crackle or explosion with smoke at the end. Potassium reacts vigorously with water. The picture shows the lilac flame as potassium reacts on the surface of the water with sparks being given off. pg 53  GCSE CHEMISTRY: GUIDANCE The Periodic Table Reactions of Group 1 metals with water: https://www.youtube.com/watch?v=6ZY6d6jrq-0 The organisation of the Periodic Table: https://www.bbc.co.uk/bitesize/guides/z6pct39/video The general word equation for a metal reacting with water is: metal + water → metal hydroxide + hydrogen This applies to all metals reacting with water (including those in Group 2). The hydrogen gas causes the fizzing and the solution remaining will be a strong alkali as the Group 1 metal hydroxides are strong alkalis. This can be shown by adding some indicator in the solution. The balanced symbol equations for the reactions of the Group 1 metals with water are: 2Li + 2H2O → 2LiOH + H2 2Na + 2H2O → 2NaOH + H2 2K + 2H2O → 2KOH + H2 State symbols are often added to these equations, for example: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Exam Tip The balanced symbol equations for Group 1 metals reacting with water are all balanced as 2:2:2:1. These are difficult equations to balance so it can be easier to remember this. The alkali metals have similar chemical properties because when they react the atoms each lose one electron to form a positive ion which is stable as it has a full outer shell of electrons. The half equation for this process for sodium is: Na → Na+ + e- As we go down the group, the outer electron in the atom is further from the nucleus so is less attracted to the nucleus. The outer electron is lost more easily from atoms further down the group so reactivity increases down the group. Solid Group 1 compounds are white solids and dissolve in water to give colourless solutions. pg 54  GCSE CHEMISTRY: GUIDANCE The Periodic Table Test yourself 1.6.2 1. When potassium reacts with water, heat is released and a colourless solution is formed. State four other observations you would make during this reaction. 2. Explain why Group 1 elements have similar chemical properties. 3. Expl