IGCSE Chemistry - Topic 1 - States of Matter PDF

IGCSE Chemistry – Topic 1 Mrs. Larisa Thomas STATES OF MATTER PROPERTIES OF SOLIDS, LIQUIDS GASES  SOLIDS Solids have a fixed volume and shape They have a high density The atoms vibrate in position but can’t change location The partic...

IGCSE Chemistry – Topic 1 Mrs. Larisa Thomas STATES OF MATTER PROPERTIES OF SOLIDS, LIQUIDS GASES  SOLIDS Solids have a fixed volume and shape They have a high density The atoms vibrate in position but can’t change location The particles are packed very closely together in a fixed and regular pattern This regular pattern is called a lattice  LIQUIDS Liquids have a fixed volume but adopt the shape of the container They are generally less dense than solids (except water) but much denser than gases The particles move and slide past each other This is why liquids adopt the shape of the container and why they can flow freely  GASES Gases don’t have a fixed volume and take up the shape of the container They have very low density There is a lot of space between the particles so gases can be compressed into a smaller volume Particles are far apart and move randomly and quickly in all directions They collide with each other and the sides of a container This is how pressure is created inside a can of gas State Solid Liquid Gas Particle Particles close together: Particles close together: Particles far apart Separation tightly packed loosely packed Arrangement of Regular pattern - lattice Randomly arranged Randomly arranged Particles Motion of Vibrate around a fixed Move or slide around each Move quickly in all directions Particles position other (straight lines) Energy of Low kinetic energy Medium kinetic energy High kinetic energy Particles Density High Medium Low Not fixed – Fluid shape Not fixed – Fluid shape Shape Fixed shape Takes shape of container Takes shape of container Not fixed – expands to fill Volume Fixed volume Fixed volume container 2D Diagram SOLID LIQUID GAS CHANGES OF STATE  MELTING Melting is when a solid changes into a liquid Requires an increase in temperature by supplying heat energy which is transformed into kinetic energy, allowing the particles to vibrate more vigorously Occurs at a specific temperature known as the melting point (m.p.) FREEZING Freezing is when a liquid changes into a solid This is the reverse of melting and occurs at the same temperature The melting and freezing point of a pure substance is the same Requires a significant decrease in temperature (loss of heat energy) and occurs at a specific temperature known as the freezing point  EVAPORATING When a liquid changes into a gas over a range of temperatures Evaporation occurs only at the surface of liquids where high energy particles can escape from the liquid's surface at low temperatures (below the b.p.) The larger the surface area and the warmer the surface of the liquid, the quicker a liquid can evaporate BOILING Boiling is when a liquid changes into a gas Heating causes bubbles of gas to form below the surface of a liquid (liquid particles escape from the surface and within the liquid) Occurs at a specific temperature known as the boiling point (b.p.)  CONDENSING When a gas changes into a liquid on cooling and it takes place over a range of temperatures When a gas is cooled its particles lose energy and when they bump into each other they lack the energy to bounce away again, instead they group together to form a liquid Changes of State  What is the state? With information about the melting and boiling points of a substance, one can identify the state the substance is in at a specific temperature If the given temperature is below the melting point (freezing point) → the substance will be a solid at that temperature If the given temperature is between the melting point and boiling point→ the substance will be a liquid at that temperature If the given temperature is above the boiling point→ the substance will be a gas at that temperature  Melting Boiling point point SOLID LIQUID GAS Between melting and Below melting point boiling point Above boiling point  The Kinetic Particle Theory The Kinetic Particle Theory states the following All matter is made up of very small particles The particles are in constant random motion → They have kinetic energy There are spaces between the particles → Intermolecular spaces There are attractive forces between the particles → Intermolecular forces  State Changes - Kinetic Theory MELTING → A solid is heated and the heat energy supplied is transformed into kinetic energy. Increased kinetic energy causes stronger vibrations until the particles have enough energy to weaken the intermolecular forces holding them in a regular arrangement. The intermolecular spaces will increase as particles break away from the lattice arrangement to form a liquid. FREEZING → As a liquid is cooled down, the kinetic energy of the particles decrease, and they start moving slower. At a certain temperature, their motion becomes slow enough for the forces of attraction to be able to hold the particles together in a regular arrangement of a solid. As the intermolecular forces become stronger, the intermolecular spaces become smaller. BOILING → A liquid is heated, and the heat energy supplied is transformed into kinetic energy. An increase in kinetic energy causes the particles to move faster and further until the particles move fast enough to overcome (break all) the intermolecular forces holding them together. The intermolecular spaces will increase as particles break away from the liquid arrangement to form a gas. CONDENSING → As a gas is cooled, the kinetic energy of the particles decrease, and they start moving slower. At a certain temperature, the gas particles will slow down enough for the attractive forces to become strong enough to hold them together in a liquid arrangement. As the intermolecular forces become stronger, the intermolecular spaces become smaller.  State Changes & Kinetic Theory - Summary When substances are heated, the particles absorb heat (thermal) energy which is converted into kinetic energy An increase in kinetic energy in a solid causes the particles to vibrate more and as the temperature increases, they vibrate so much that the solid expands until the structure breaks and the solid melts On further heating, the particles in the now liquid substance also absorbs heat energy, which is converted into kinetic energy, causing the particles to move more and faster The liquid expands more and some particles at the surface gain enough energy to overcome the intermolecular forces and evaporate When the boiling point temperature is reached, all the particles gain enough energy to escape, and the liquids boils  HEATING AND COOLING CURVES  State Changes on Graphs Changes in state can be shown on graphs called heating curves and cooling curves. These curves show how changes in temperature affect changes of state.  A heating curve shows the change of state of a Heating Curve substance from solid to gas when it is heated  Interpreting a Heating Curve Heating a solid results in its temperature rising over the time of heating but the graph shows two periods during which the temperature remains constant The temperature is constant during the phase changes → melting and vaporisation In the regions where the temperature rises: The heat energy added is transformed into kinetic energy so temperature increases Increased kinetic energy causes the particles to move faster and interact less strongly The intermolecular spaces increase as the particles begin to move apart In the regions where the temperature is constant: The heat energy added is used to overcome the intermolecular forces The heat energy causes changes in potential energy NOT kinetic energy This results in the temperature staying constant until the phase change is complete  Plateau = Phase change = Potential energy change  A cooling curve shows the change of state of a Cooling Curve substance from gas to solid when it is cooled  Interpreting a Cooling Curve Cooling a gas results in its temperature falling over the time of cooling The temperature is constant during the two phase changes → condensing and freezing In the regions where the temperature falls: Kinetic energy is transformed into heat energy that is removed so temperature decreases Decreased kinetic energy causes the particles to move slower and interact stronger The intermolecular spaces decrease as the particles begin to move closer together In the regions where the temperature is constant: The heat energy removed comes from energy released when forming new intermolecular forces The heat energy comes from changes in potential energy NOT kinetic energy The temperature stays constant until the phase change is complete  Pure VS Impure Substances Pure Substance → A substance that consists of only one type of element or compound Impure Substance → A substance that consists of more than one type of element and/or compound not chemically bonded Mixtures (pure substances physically mixed with impurities) are impure substances Impurities causes the melting point of an impure substance to be lower than the pure substance and causes it to melt over a range of temperatures Impurities causes the boiling point of an impure substance to be higher than the pure substance and causes it to boil over a range of temperatures Pure substances have specific and fixed melting and boiling points Phase changes on a heating or cooling curve are horizontal/flat lines Impure substances melt and boil over a range of temperatures Phase changes on a heating or cooling curve are slope lines with a gradient Pure Substances Heating Curve A pure substance boils at a specific and constant temperature A pure substance melts at a specific and constant temperature Impure Substances Heating Curve An impure substance boils at a higher temperature and over a range of temperatures An impure substance melts at a lower temperature and over a range of temperatures VOLUME OF GASES  Effect of Temperature & Pressure on the VOLUME of a Gas Inversely Changing the external pressure on a sample of gas Proportional An increase in pressure produces a decrease in volume → Gas is compressed A decrease in pressure produces an increase in volume → Gas expands Directly Changing the temperature of a sample of gas Proportional An increase in temperature produces an increase in volume → Gas expands A decrease in temperature produces a decrease in volume → Gas is compressed The large intermolecular spaces in gases explain why the volume is easily changed by changes in temperature and pressure Effect of Temperature on the Volume of a Gas Kinetic Theory Increase in temperature → The kinetic energy of the gas particles increase; they move faster and there is less chance of interaction between them as the intermolecular forces become almost negligent. They can move further apart to occupy a greater volume. Decrease in temperature → The kinetic energy of the gas particles decrease; they move slower, and they are more likely to interact with each other as the intermolecular forces have a greater effect. They will move closer together to occupy a smaller volume.  Effect of Pressure on the Volume of a Gas Kinetic Theory Increase in pressure→ The gas particles are pushed closer together and are more likely to interact with each other as the intermolecular forces have a greater effect. They will move closer together to occupy a smaller space/volume. Decrease in pressure→ The gas particles are not pushed together and are less likely to interact with each other as the intermolecular forces will have less of an effect. They will move further apart and occupy a greater space/volume.  Gas Pressure Gas particles are in constant and random motion The pressure that a gas creates inside a closed container is produced by the gas particles hitting the inside walls of the container  Effect of Temperature and Volume on Gas Pressure IINCREASE in TEMPERATURE → The heat energy supplied is transformed into kinetic energy. This increases the kinetic energy of the gas particles, so they move faster and collide with the walls of the container more frequently. The pressure will increase. DECREASE in VOLUME → If the container is made smaller, the same amount of gas particles will have less space available to move in and will collide with the walls of the container more frequently. The pressure will increase. Decreasing the volume of a container increases the gas pressure DIFFUSION  Diffusion The movement of particles from an area of higher concentration to an area of lower concentration, down a concentration gradient, until equilibrium is reached Equilibrium means that eventually the concentration of particles will be equal as they spread out evenly to occupy all the available space This is the process by which different gases or liquids mix and is due to the random motion of their particles Diffusion can only happen in fluids (liquids and gases) because they have large enough intermolecular spaces for the particles to move around and spread out Diffusion cannot happen in solids because the intermolecular spaces are very small and the particles only vibrate, they do not move around to spread out evenly Diffusion of potassium manganate(VII), KMnO4 , in water. After a few hours, the concentration of KMnO4 is the same throughout the solution  Diffusion Rate & Molecular Mass Diffusion occurs much faster in gases than in liquids as gaseous particles move much quicker than liquid particles At the same temperature, different gases do not diffuse at the same rate This is due to the difference in their relative molecular masses Particles with a lower relative molecular mass is lighter, can move faster and further and therefore will diffuse at a faster rate Particles with a higher relative molecular mass is heavier, they move slower and therefore will diffuse at a slower rate  This can be demonstrated in the reaction between ammonia (NH3) and hydrogen chloride gas (HCl) inside a long glass tube Where the two gases meet, a white smoke of ammonium chloride (NH4Cl) forms This does not occur in the middle of the tube, but much closer to the end with the hydrogen chloride Hydrogen chloride has a relative molecular mass of 36.5 and ammonia of 17 The ammonia molecules are lighter, move faster and thus diffuse faster  Diffusion Rate & Temperature Diffusion is a passive process, which means that it happens on its own and no energy input is required Particles of the same substance has the same molecular mass but will diffuse at different rates if the temperature differs For the same substance, the rate of diffusion is faster at a higher temperature as the particles will have more kinetic energy and move faster For the same substance, the rate of diffusion is slower at a lower temperature as the particles will have less kinetic energy and move slower Topic Summary ATOMS, ELEMENTS AND COMPOUNDS IGCSE Chemistry – Topic 2 Mrs. Larisa Thomas  ELEMENTS, COMPOUNDS AND MIXTURES  Atoms All matter is made of atoms. Very long ago, a Greek philosopher called Democritus said that all matter is made up of tiny pieces. He suggested that if you take a substance and keep cutting it into smaller pieces, you will end up with a piece that cannot be cut anymore or cannot get any smaller. He called these smallest pieces of matter atoms, because the word “atom” means “cannot be divided”. Atom → The smallest particle of a substance, that cannot be broken down chemically.  Substances Atoms are the building blocks of substances. All substances can be classified into one of these three types: Elements Compounds Mixtures Element A pure substance made up of only one type of atom, each containing the same number of protons, and it cannot be split into simpler substances There are 118 elements found on the Periodic Table E.g. hydrogen and magnesium Compound A pure substance made up of two or more elements chemically combined There is an unlimited number of compounds Compounds cannot be separated into their elements by physical means E.g. copper(II) sulfate (CuSO4), calcium carbonate (CaCO3), carbon dioxide (CO2) Mixture An impure substance made up of two or more substances (elements and/or compounds) that are not chemically combined Mixtures can be separated by physical methods such as filtration or evaporation E.g. sand and water, oil and water, sulfur powder and iron filings  Substance Explanation Particle Model Particle Model Consists of only one type of Element atom (may be diatomic) Consists of two or more Compound different elements (atoms) chemically bonded Consists of two or more Mixture of elements different elements NOT chemically bonded Consists of two or more Mixture of different compounds NOT compounds chemically bonded Consists of two or more Mixture of elements different elements and and compounds compounds NOT chemically bonded  ATOMS & ATOMIC STRUCTURE  Atoms are the building blocks of matter Each atom is made of subatomic particles General Atomic called protons, neutrons and electrons The protons and neutrons are located at Structure the centre of the atom, called the nucleus The electrons move very fast around the nucleus in orbital paths called shells The mass of the electron is so small it can be ignored; which means the mass of an atom is contained within the nucleus where the protons and neutrons are located The number of subatomic particles and shells in an atom varies for every element  Atomic Structure and Mass Atomic structure can thus be described as: A central nucleus containing protons and neutrons surrounded by electrons in shells Atoms are so small that we cannot compare their masses in conventional units like grams A unit called the relative atomic mass is used This is also used to measure the mass of the subatomic particles  Subatomic Particles Subatomic particles are often compared and identified by means of their relative mass measured in relative atomic mass units and electric charge as either having a positive, negative or neutral (no) charge Particle Relative mass Charge Proton 1 1+ Neutron 1 0 1 Electron OR 0 1- 1840  Atomic Notation The atomic number and mass number of an element can be shown using atomic notation Proton/Atomic Number (Z) The number of protons in the nucleus of an atom (symbol Z) It determines the position of the element on the Periodic Table Nucleon/Mass Number (A) The number of protons and neutrons in the nucleus of an atom (symbol A) Protons and neutrons can collectively be called nucleons The Periodic Table shows elements with their atomic/proton number at the top and relative atomic mass at the bottom - there is a difference between relative atomic mass and mass number, but for your exam, you can use the relative atomic mass as the mass number (except for chlorine)  Calculating the number of protons, neutrons and electrons Protons The atomic number of an atom or ion determines which element it is All atoms or ions of the same element have the same number of protons E.g., Li has an atomic number of 3, so all atoms/ions of Li will have 3 protons Number of protons = Atomic (proton) number The number of protons of an unknown element can be calculated by using its mass number and number of neutrons: Number of protons = mass number – number of neutrons Electrons for Atoms Atoms have an overall neutral charge because the positive charge on the protons and negative charge on the electrons balance (cancel out) This means that an atom has the same number of protons and electrons Number of electrons = Number of protons = Atomic (proton) number Neutrons The mass and atomic numbers are used to find the number of neutrons in atoms or ions The mass number shows the total number of protons and neutrons, so to get the number of neutrons only, you need to subtract the number of protons Number of neutrons = mass number – atomic (proton) number Example Determine the number of protons, electrons and neutrons in an atom of unknown element X with atomic number 29 and mass number 63. Protons: Number of protons = the atomic number 63 Number of protons = 29 Electrons: Atoms are neutral Number of electrons = number of protons 29 𝑋 Number of electrons = 29 Neutrons: Number of neutrons = mass number – number of protons Number of neutrons = 63 – 29 Number of neutrons = 34 Example Using the periodic table, calculate the number of protons, neutrons and electrons in an atom of nitrogen Atomic 7 number Number of protons = atomic number = 7 N nitrogen Number of electrons = number of protons = 7 Atomic 14 mass Number of neutrons = atomic mass – atomic number = 14 - 7 =7 Example An atom of a specific element contains 9 protons, 9 electrons and 10 neutrons. An atom of which 9 element is this? F fluorine The atomic number determines which element it 19 is. The atom has 9 protons, so the element is fluorine because fluorine has an atomic number of 9 on the periodic table. ELECTRONIC STRUCTURE  Electron Shell Diagrams Drawing the electronic Electrons orbit the nucleus in shells (energy levels) structure Electrons start filling the shell closest to the nucleus When a shell is full, the additional electrons start filling the next shell Different shells can hold a different number of electrons before being full: The first shell can hold 2 electrons The second shell can hold 8 electrons The third shell can hold 8 electrons (simplified for this course) The outermost shell of an atom is called the valence shell, and an atom is much more stable if it can manage to completely fill this shell with electrons The electrons on the outermost shell are called valency electrons Electrons can be represented as either dots ( ) or crosses (x) Electron Shell Diagrams & The Periodic Table The electron shell diagram for an atom of any element can be drawn by using information from the Periodic Table The Periodic Table is arranged into 7 periods (→) and 8 groups (↓) The period number that an element is found in, determines how many electron shells an atom for that element will have The group number that an element is found in, determines how many valency electrons an atom for that element will have Once you have filled the inner shells with their maximum number of electrons, draw the correct number of electrons on the outermost shell To double check, make sure the total number of electrons is the same as the atomic (proton) number for the element  Example: Group IV (4) elements have Shows the number atoms with 4 electrons in the outer of valency electrons shell & Group VI (6) elements have atoms with 6 valency electrons Shows the number of I II GROUPS III IV V VI VII O electron shells 1 2 PERIODS 3 4 Example: 5 Elements in period 2 have 2 electron 6 shells & elements in period 3 have 3 7 electron shells  Writing the Electronic Configuration electronic structure The arrangement of electrons in shells can also be explained using numbers instead of electron shell diagrams Thisspecial notation called the electronic configuration/electronic structure/electronic distribution The number of electrons in each shell, starting from the first shell and moving out, can be written down, separated by commas Example → Carbon has 6 electrons, 2 in the first shell and 4 in the second shell Its electronic configuration is 2,4 Electron Configuration & The Periodic Table The number of notations in the electronic configuration shows the number of electron shells the atom has and the period that element is in The last notation shows the number of valency electrons the atom has, showing the group that element is in Example Draw and write the electronic structure of magnesium. Mg is in Period 3, so it has 3 electrons shells The first shell is filled with two electrons and the second with eight Mg is in Group 2, so it has 2 valency electrons in the third shell The written form of this The atomic number of Mg is 12 so electronic structure is 2,8,2 Mg has 12 electrons in total  IONS & IONIC STRUCTURE  Ions Ion → An electrically charged atom or group of atoms formed by the loss or gain of electrons. This loss or gain of electrons are to obtain a full outer shell of electrons The electronic structure of ions will be the same as that of a noble gas Negative ions are called anions and form when atoms gain electrons They have more electrons (-) than protons (+) Non-metals gain electrons from other atoms to become anions Positive ions are called cations and form when atoms lose electrons They have more protons (+) than electrons (-) Metals lose electrons to other atoms to become cations  𝑻𝒉𝒆 𝒄𝒉𝒂𝒓𝒈𝒆 𝒐𝒏 𝒂𝒏 𝒊𝒐𝒏 𝒊𝒔 Charge on Ions 𝒊𝒏𝒅𝒊𝒄𝒂𝒕𝒆𝒅 𝒂𝒔 𝒂 𝒔𝒖𝒑𝒆𝒓𝒔𝒄𝒓𝒊𝒑𝒕 𝒂𝒇𝒕𝒆𝒓 𝒕𝒉𝒆 𝒔𝒚𝒎𝒃𝒐𝒍 𝒆. 𝒈. 𝑴𝒈𝟐+ Atoms have no charge because they have an equal number of protons (+) and electrons (-), while ions have a positive or negative charge because they do not have an equal number of protons and electrons When atoms gain electrons, they form anions with more electrons than protons The size of the charge depends on how many electrons are gained Oxygen (Group 6) has 6 electrons on the outer shell and will gain 2 electrons to obtain a full outer shell of 8 electrons → An oxygen ion will have a 2- charge When atoms lose electrons, they form cations with less electrons than protons The size of the charge depends on how many electrons are lost Magnesium (Group 2) will lose its outer shell with 2 electrons for its previous full shell to become its outer shell → A magnesium ion will have a 2+ charge  Ionic Charge & The Periodic Table There is a link between group numbers and the charge on ions as the number of valency electrons an atom has determines the number of electrons it will lose or gain to form an ion 4+ 1+ 2+ 3+ 4- 3- 2- 1- I II GROUPS III IV V VI VII O  Electronic Configuration of Ions Electronic configurations can also be written for ions A sodium atom has 11 electrons: 2 in the first shell, 8 in the second shell and 1 in the third shell Electronic configuration → Na 2,8,1 A sodium ion has lost its one outer electron and outer shell, therefore has 10 electrons: 2 in the first shell and 8 in the second shell Electronic configuration → Na+ 2,8 Read carefully whether you are asked to give the electronic structure of atoms or ions in the exam and remember to show the charge if you are working with ions  FORMATION OF A CATION CATions are The mass and atomic numbers don’t 𝟐𝟑 change as the number of protons and 𝟐𝟑 + PAWsitive 𝟏𝟏𝑵𝒂 neutrons don’t change. The + charge 𝟏𝟏 𝑵𝒂 shows the loss an electron!  FORMATION OF AN ANION The mass and atomic numbers don’t 𝟑𝟒 change as the number of protons and 𝟑𝟒 − 𝟏𝟕𝑪𝒍 neutrons don’t change. The - charge 𝟏𝟕𝑪𝒍 shows the gain of an electron! The number protons, neutrons and electrons Protons and Neutrons To get the number of protons and neutrons for an ion is the same as for an atom Number of protons = Atomic (proton) number Number of neutrons = mass number – atomic (proton) number Electrons for Ions Ions have an overall positive or negative charge because ions do not have the same number of protons and electrons Number of electrons in anion (-) = number of protons + size of charge Number of electrons in cation (+) = number of protons - size of charge  Polyatomic Ions A group of bonded atoms that carry a net charge because the total number of electrons in the group is not equal to the total number of protons Polyatomic ion Formula Charge Formula and Charge Ammonium NH4 + NH4+ Hydrogen carbonate HCO3 - HCO3- Hydroxide OH - OH- Nitrate NO3 - NO3- Carbonate CO3 2- CO32- Sulfate SO4 2- SO42- Phosphate PO4 3- PO43- ISOTOPES  Isotopes Different atoms of the same element that contain the same number of protons but a different number of neutrons Due to the number of protons being the same but the number of neutrons being different, isotopes will have the same atomic number but different mass numbers The symbol for an isotope is the chemical symbol (or word) followed by a dash and then the mass number C-14 or carbon-14, also written as 146𝐶 or 14𝐶 is the isotope of carbon with a mass number of 14, 6 protons and 8 neutrons Isotopes of Hydrogen  Isotopes Share Chemical Properties Isotopes of the same element have the same chemical properties because they have the same number of electrons / electronic configuration The number of electrons, specifically those on the outer shell, determines the chemistry of an atom The difference between isotopes is the number of neutrons, which are neutral particles within the nucleus that add mass only The difference in mass between isotopes affects the physical properties, such as density, boiling point and melting point Isotopes are identical in appearance and reaction, so a sample of C-14 would look and react no different to C-12  Relative Atomic Mass Atoms are so tiny that we can’t measure their masses in units like grams, so a unit called the relative atomic mass (Ar) is used One relative atomic mass unit is 1/12th the mass of a carbon-12 atom All other elements are measured in comparison to the mass of a C-12 atom and since these are ratios, the relative atomic mass has no units E.g., hydrogen has a relative atomic mass of 1, meaning that 12 atoms of hydrogen would have the same mass as 1 atom of carbon The Ar for each element appears on the Periodic Table and is an average mass of all the isotopes of that element, rounded off to a whole number  Calculating Relative Atomic Mass As the relative atomic mass is the average mass of all the isotopes of an element, it takes into account how many of each isotope is present The relative atomic mass of an element is calculated from the mass number and relative abundances (%) of all its isotopes To calculate the relative atomic mass, the equation below is used where the top line is the sum of all the different isotopes of a particular element 𝜮 (𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 × 𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝒎𝒂𝒔𝒔 𝒏𝒖𝒎𝒃𝒆𝒓) 𝑨𝒓 = 𝟏𝟎𝟎 Example IONIC BONDS  ionic compound Ionic Bonding Ionic Bond → A strong electrostatic attraction between oppositely charged ions electrostatic attraction (ionic bond) Ionic compounds form when metals react with non-metals The metal atoms lose their valency electrons to form cations (+), and the non-metal atoms gain these electrons to form anions (-) Electrons are transferred from metal atoms to non-metal atoms The positive and negative ions are held together by strong electrostatic forces of attraction between opposite charges This force is known as an ionic bond, and they hold ionic compounds together  These diagrams show the arrangement of the electrons in an ionic compound Dot-and-cross Electrons are shown as dots and Diagrams crosses: In a dot and cross diagram for ionic compounds: The electrons for one element/atom is shown with dots and the electrons for a different element is shown with crosses The charge of the ion is spread evenly which is shown by using large square brackets The charge on each ion is written at the top right-hand corner  Ionic Bonds between Group 1 and Group 7 Elements Group 1 metals lose their one valency electron to obtain a full outer shell They form positive ions with a charge of 1+ Group 7 non-metals gain this one electron to obtain a full outer shell They form negative ions with a charge of 1– One electron is transferred from the outer shell of the Group 1 metal atom to the outer shell of the Group 7 non-metal atom to form ions The oppositely charged Group 1 cation and Group 7 anion is then attracted to one another and held together by electrostatic forces  Example: Sodium reacts with chlorine to form sodium chloride One valency electron is transferred from the sodium atom to the chlorine atom This forms a sodium cation with a 1+ charge and a chloride anion with 1- charge The oppositely charged ions are held together by electrostatic forces of attraction DOT-AND-CROSS DIAGRAM OF SODIUM CHLORIDE COVALENT BONDS  covalent compound: molecule Covalent Bonding Covalent Bond → A bond formed when a pair of electrons is shared between two atoms, leading to noble gas electronic configurations covalent bond Covalent compounds are formed when pairs of electrons are shared between atoms for each atom obtain a full outer shell of electrons Only non-metal elements participate in covalent bonding Covalent substances may consist of small molecules or giant molecules When two or more atoms are covalently bonded together, we describe them as ‘molecules’  These diagrams show the arrangement of Dot-and-cross the electrons in a molecule Electrons are shown as dots and crosses: Diagrams In a dot and cross diagram for covalent compounds: The electrons for one atom/element is shown with dots and the electrons for the other atom is shown with crosses The electron shells of each atom in the molecule overlap and the shared electrons are shown in this area The electron shell for each atom must be full, counting the shared electrons for each atom  FORMATION OF A COVALENT BOND SIMPLE COVALENT MOLECULES HYDROGEN H2 CHLORINE Cl2 WATER H2O METHANE CH4 AMMONIA NH3 HYDROGEN CHLORIDE HCl CHEMISTRY OF THE ENVIRONMENT IGCSE Chemistry – Topic 10 Mrs. L Thomas WATER  Chemical Tests for Water Using Anhydrous Cobalt Chloride Blue anhydrous cobalt (II) chloride turns pink on the addition of water This test is usually done using cobalt chloride paper anhydrous cobalt (II) chloride + water ⇌ hydrated cobalt (II) chloride CoCl2(s) + 6 H2O(l) ⇌ CoCl2∙6H2O(s) Anhydrous → Hydrated Blue → Pink  Chemical Tests for Water Using Anhydrous Copper Sulfate White anhydrous copper (II) sulfate turns blue on the addition of water This test is usually done using copper sulfate powder/crystals anhydrous copper (II) sulfate + water ⇌ hydrated copper (II) sulfate CuSO4(s) + 5 H2O(l) ⇌ CuSO4∙ 5H2O(s) Anhydrous → Hydrated White → Blue  Purity of Water – Physical Tests Testing for the purity of water can be done with physical tests Pure substances melt and boil at specific and sharp temperatures Water has a boiling point of 100 °C and a melting point of 0 °C Mixtures/Impure substances have a range of melting and boiling points as they consist of different substances that melt or boil at different temperatures Melting and boiling points data can be used to determine the purity of water Impurities tend to increase the boiling point of water, so impure water will start boiling at temperatures above 100 °C Impurities tend to decrease the melting point of water, so impure water will start melting at temperatures below 0 °C  Distilled Water Distilled water is water that has been heated to form a vapour, and then condensed back to a liquid It contains very few chemical impurities Distilled water is used in practical chemistry because of its high purity Tap water contains more impurities which could interfere with chemical reactions so is typically not used in practicals  Water from Natural Sources We use water in many aspects of our everyday life: Domestic uses: drinking, cooking, gardening and general sanitation Agricultural uses: drink for animals and watering crops Industrial uses: as a solvent and coolant and to generate electricity Natural sources of water include lakes, rivers and underground water sources (groundwater) A rock that stores water is known as an aquifer  Substances in Water from Natural Sources Water from natural sources may contain a variety of different substances, including: ❖ Dissolved oxygen ❖ Metal compounds Some of these substances ❖ Plastics are naturally occurring but many are a direct ❖ Sewage result of human activities ❖ Harmful microbes ❖ Nitrates from fertilisers ❖ Phosphates from fertilisers and detergents Many of these substances enter water sources when rain falls and washes them into lakes, rivers or groundwater  Beneficial Substances in Water Some of the substances found in natural water sources are beneficial and others may have harmful effects Beneficial substances include: ❖ Dissolved oxygen - essential for aquatic life ❖ Metal compounds - some provide essential minerals which are necessary for life, such as calcium and magnesium  Harmful Substances in Water Potentially harmful substances include: ❖ Metal compounds - some are toxic like aluminium and lead ❖ Some plastics - these may harm aquatic life in many ways, e.g., getting trapped in plastic waste, dying of starvation as their stomach fill with plastic ❖ Sewage - contains harmful microbes which can cause disease ❖ Nitrate & phosphates - these can promote the growth of aquatic plants which leads to the deoxygenation of water. Ultimately, this can cause damage to aquatic life in a process called eutrophication  Water Treatment Untreated water is taken from rivers, reservoirs or underground water sources (groundwater) and is treated to make it safe for use as domestic water Untreated water contains soluble and insoluble impurities Insoluble impurities: soil, pieces of plants and other organic matter Soluble impurities: dissolved calcium, metallic compounds and inorganic pollutants Three main parts in the treatment of domestic water Sedimentation & Filtration → Removes solids Use of carbon → Removes tastes and odours Chlorination → Kills microbes  Water Treatment Process Water is pumped into sedimentation tanks and is allowed to stand for a few hours In a process called sedimentation; mud, sand and other particles will fall to the bottom of the tank due to gravity and form a layer of sediment Filtration is used to remove smaller particles by passing the water through layers of sand and gravel filters that trap solid particles Water can also be passed through carbon (in the form of charcoal) to remove tastes and odours Bacteria and other unwanted microorganisms are too small to be trapped by the filters. Chlorination, the careful addition of chlorine to the water supply, is used to kill them Cholera and typhoid are examples of bacterial diseases which can arise from the consumption of untreated water SUMMARY FERTILISERS  Fertilisers Ammonium salts and nitrates are commonly used as fertilisers NPK fertilisers provide the elements nitrogen, phosphorus and potassium for improved plant growth Different fertilisers contain different amounts of fertiliser compounds, so each contains different proportions of nitrogen, potassium and phosphorous  N,P,K Fertilisers Fertilisers that contain nitrogen (N), potassium (K) and phosphorus (P) Nitrogen makes chlorophyll and protein and promotes healthy leaves Potassium promotes growth and healthy fruit and flowers Phosphorus promotes healthy roots Fertiliser compounds contain the following water-soluble ions: Ammonium ions, NH4+, and nitrate ions, NO3-, are sources of soluble nitrogen Phosphate ions, PO43-, are a source of soluble phosphorus Most potassium compounds dissolve in water to produce potassium ions, K+ Common fertiliser compounds include: Ammonium nitrate, NH4NO3 Ammonium phosphate, (NH4)3PO4 Potassium sulfate, K2SO4 AIR QUALITY AND CLIMATE  The Composition of Air The chart shows the approximate percentages by volume of the Mixture of noble gases and carbon dioxide main gases in clean, dry air 78% Nitrogen 21% Oxygen 1% Mixture of Noble Gases and Carbon Dioxide  Uses of Air The gases available in the air have many important applications The noble gases are used in many applications, e.g., helium is used to fill balloons, argon is used in tungsten light bulbs, krypton is used in lasers for eye surgery Oxygen is used in steel making, welding and breathing apparatus Nitrogen is used in food packaging, the production of ammonia and in the production of silicon chips Oxygen and nitrogen are separated from the air by fractional distillation  Air Pollution In addition to the gases present naturally in our atmosphere, other gases are present due to human activities and are classed as air pollutants CARBON DIOXIDE Sources: complete combustion of carbon-containing fuels such as fossil fuels, e.g., the complete combustion of methane: CH4 + O2 → CO2 + 2H2O Adverse effects: higher levels of carbon dioxide can increase global warming, which leads to climate change  CARBON MONOXIDE Sources: incomplete combustion of carbon-containing fuels such as fossil fuels, e.g., incomplete combustion of gasoline: C8H18 + 9 O2 → 5 CO + 2 CO2 + 9 H2O Adverse effects: toxic gas as combining with haemoglobin in the blood and prevents it from carrying oxygen PARTICULATES Sources: incomplete combustion of carbon-containing fuels such as fossil fuels can also produce particulates of carbon (soot), e.g., the incomplete combustion of methane can produce CO and C: 2 CH4 + 3 O2→ 2 CO + 4 H2O CH4 + O2→ C + 2 H2O Adverse effects: increased risk of respiratory problems and cancer  METHANE Sources: waste gases from digestion in animals, decomposition of vegetation and bacterial action in swamps, rice paddy fields and landfill sites Adverse effects: higher levels of methane leading to increase global warming and climate change OXIDES OF NITROGEN Sources: reaction of nitrogen with oxygen in the presence of high temperatures, e.g., in car engines, high-temperature furnaces and lightning Adverse effects: Produces photochemical smog Dissolves in rain to form acid rain which causes damage to aquatic organisms and corrosion to metal structures/buildings/statues made of carbonate rocks Causes respiratory problems and irritates the lungs, throat and eyes  SULFUR DIOXIDE Sources: combustion of fossil fuels containing sulfur compounds. Power stations are a major source. Adverse effects: dissolves in rain to form acid rain with similar effects as the acid rain caused by oxides of nitrogen.  How Greenhouse Gases Cause Global Warming The Sun emits energy in the form of radiation that enters the Earth’s atmosphere Most thermal energy is absorbed and re-emitted back from the Earth’s surface → The thermal energy passes through the atmosphere, and some is emitted straight into space Some thermal energy is absorbed by greenhouse gases such as carbon dioxide and methane and is then reflected and emitted in all directions This reduces thermal energy loss to space and traps it within the Earth’s atmosphere This process is known as the greenhouse effect As the concentration of greenhouse gases in the atmosphere increases due to human activity, more thermal energy is trapped within the Earth's atmosphere causing the Earth’s average temperature to rise (global warming) This process is called the enhanced greenhouse effect Consequences of Global Warming Climate change due to the increase in Earth’s temperature This can lead to a loss of habitat Water levels will rise as glaciers melt, causing flooding in low-lying countries Extinction of species due to the destruction of natural habitats Migration of species as they move to areas that are more habitable Spread of diseases caused by warmer climate  Reducing the Effects of Environmental Issues CLIMATE CHANGE The production of greenhouse gases needs to be reduced drastically to avoid or at least slow climate change CO2 emissions can be reduced by increasing the use of hydrogen and renewable energy such as solar or wind energy to decrease the use of fossil fuels Reduce livestock farming to decrease the methane emissions produced from digestion in animals Planting more trees to remove more carbon dioxide from the atmosphere  ACID RAIN – SULFUR DIOXIDE Emissions of sulfur dioxide can be reduced by either: Using low-sulfur fuels Flue gas desulfurisation → this involves reacting the sulfur dioxide emitted from burning fuels with calcium oxide to remove it from the gas ACID RAIN - OXIDES OF NITROGEN NO and NO2 are formed when nitrogen and oxygen react under high pressure and temperatures in internal combustion engines and blast furnaces Exhaust gases also contain unburned hydrocarbons and carbon monoxide Cars are fitted with catalytic converters which form a part of their exhaust systems and act to render these vehicle exhaust gases harmless  Catalytic Converters They contain a series of transition metal catalysts like platinum and rhodium in a honeycomb structure which increases the reaction surface area A series of redox reactions occurs which neutralises the pollutant gases Carbon monoxide is oxidised to carbon dioxide: 2 CO + O2 → 2 CO2 Oxides of nitrogen are reduced to nitrogen gas: 2 NO → N2 + O2 2 NO2 → N2 + 2 O2 A single reaction can summarise the reactions within a catalytic convertor: 2 NO + 2 CO → N2 + 2 CO2  Catalytic Converters  Photosynthesis Photosynthesis is the reaction between water and carbon dioxide to produce glucose and oxygen in the presence of chlorophyll, using energy from light Chlorophyll (in chloroplast) and energy from light are required for this reaction The word equation for photosynthesis is: Reactants carbon dioxide water The balanced symbol equation for photosynthesis is: Products glucose oxygen Summary