T1: Sub-atomic particles • Atoms are made from smaller particles called subatomic particles. • There are three types we need to know about, summarised below. T1: Mendeleev • Arranged elements by increasing atomic mass but …. • He broke this rule and left some gaps if an element’s properties weren’t similar to the one above it. • He thought the gaps were for elements that hadn’t been discovered yet and predicted their properties. • When they were discovered, the properties matched the predictions PERIODS….increasing atomic mass, differing properties G R O U P S … … s i m i l a r p r o p e r t i e s Element Type = non-metal = metal Particl e Relativ e charge Relative mass Found? Proton 1 Positive , +1 In nucleus Neutron 1 Neutral, 0 In nucleus Electro n Neglibl e () Negative , -1 In shells orbiting nucleus T1: Reading the Periodic Table •Note: on some periodic tables, they are the wrong way up, just remember that the smaller number is the proton number. Relative Atomic Mass (aka nucleon number): The total number of protons and neutrons added together. Atomic number (aka proton number): The number of protons or electrons. T1: What’s in my atom? Protons = atomic number Electrons = atomic number Neutrons = relative atomic mass . – atomic number Atomic number = 9 Relative Atomic mass = 19 Protons = 9 Electrons = 9 Neutrons = 19-9 = 10 T1: Atoms and Elements • Element = substance containing only one type of atom. • Protons and electrons: same for every atom of an element…it is the number of protons that decides the element. • Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes T1: Relative Atomic Mass •This is the mass of an element relative to 1/12 th the mass of 12 C. • Element: substance containing only one type of atom. • Protons and electrons: same for every atom of an element…it is the number of protons that decides the element. • Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes. T1: Isotopes (HT) • Versions of an element with same atomic number but different atomic mass. • Number of protons is the same, but number of neutrons is different. • Relative Atomic Mass is average of the masses of the isotopes, weighted by their relative abundance • For example, Neon has three isotopes • Relative atomic mass of Neon = • This is why some atoms have a relative atomic mass with a decimal point. T1: Electron Configuration • Electrons orbit the nucleus in shells. • First shell holds two electrons • Second and third shell hold 8 electrons • Note: the third shell can actually hold more, but we won’t worry about this until A-level. Example: Silicon Atomic number is 14, so it has 14 electrons. You build up electrons from the first shell outwards, so in this case: - First shell has 2 - Second shell has 8 - Third shell has 4 This can be written as: 2.8.4; or drawn as: Neon Isotope Mass Relative Abundance (%) 20 90.5 21 0.3 22 9.2 Note: Si is in period three and group four of the periodic table; it also has three electron shells and four electrons in the outer shell – this is no coincidence! T2: Forming Ions Cations are positive (cat…pussitive!) ions They are formed when atoms lose electrons. Metals form cations by losing the electrons in their outer shells In the example, aluminium loses its three outer-shell electrons to become Al 3+ …each lost electrons cause 1 ‘+’ charge. Anions are negative ions They are formed when atoms gain electrons. Non-metals form anions by filling their outer shells. Name ends with ‘-ide’ to show it is a negative ion, In the example, oxygen gains two outer- shell electrons to become O 2- , giving it 8 electrons in its outer shell. 3+ T2: Making Ionic Compounds • An ionic bond is the attraction between a positive and a negative ion. • The overall number of positive and negative charges must cancel out. • Form between a metal and a non- metal • Ionic compounds do not form molecules Example 1: Magnesium reacting with chlorine. • Anion: Cl forms Cl - ions • Cation: Mg forms Mg 2+ ions • Formula = MgCl 2 • Why: two Cl - gives a 2- charge to balance 2+ from Mg 2+ . • Name: magnesium chloride Example 2: aluminium reacting with oxygen. • Anion: O forms O 2- ions • Cation: Al forms Al 3+ ions • Formula = Al 2 O 3 • Why: Two Al 3+ gives a 6+ charge, three O 2- gives a 6- charge. • Name: aluminium oxide T2: Ionic Structures (HT) •A repeating 3D lattice of positive and negative ions. •Strong electrosta tic bonds between ions. T2: Precipitates and Precipitation •When an insoluble salt is formed from the reaction of two soluble salts. •Goes cloudy as small particles of solid are made. • Predicting precipitates: simply choose a combination of soluble salts where you tell that if the ions swapped over you would get an insoluble salt: use the solubility table for help. • Example: T2: Common Ions • You should try to memorise the ions formed by various species: • There are also some ‘ions made of more than one atom with an overall charge: • Hydroxide: OH - • Nitrate: NO 3 - • Sulphate, SO 4 2- • Carbonate, CO 3 2- • Ammonium, NH 4 + Group Electrons in outer shell Ion forme d Examples 1 1 + Li + , Na + , K + 2 2 2+ Be 2+ , Mg 2+ , Ca 2+ 3 6 2- O 2- , S 2- 4 7 - F - , Cl - , Br - , I - T2: Solubility • Soluble: a compound dissolves in a given liquid. • Insoluble: a compound does not dissolve. Soluble in water In soluble in water All sodium, potassium, ammonium salts All nitrates Most chlorides Except: silver and lead chlorides Most sulfates Except: lead, barium and calcium sulfates. Except: sodium, potassium and ammonium carbonates Most carbonates Except: sodium, potassium and ammonium hydroxides Most hydroxides T2: Properties of Ionic Compounds • Melting point: High due to strong bonds between ions. • Boiling point: Higher, due to strong bond between ions. • Solid: do not conduct electricity • Molten (liquid): do conduct electricity • Dissolved (aqueous): do conduct electricity Why? (HT) Electrical Conductivity • Electricity is conducted when there are charged particles that are free to move. • Solid: there are charged particles (the ions), but they are not free to move, so they do not conduct. • Liquid/Aqueous: the ions are now free to move, so they do conduct High Melting/Boiling Points • Ionic bonds (attraction between positive and negative ions) are very strong. • Melting and boiling require these bonds to be broken. • This takes lots of (heat) energy. T2: Making Insoluble Salts 1. React solutions of (the right) two soluble salts together. 2. Filter the mixture to collect the precipitate. 3. Rinse the filter residue with distilled water to remove impurities. 4. Allow the residue to dry. T2: Barium Meals • A patient is given a drink containing barium sulfate. • This can show up on a x-ray, helping doctors to investigate the digestive system. T2: Flame tests 1. Clean a metal loop in acid 2. Did loop in a metal salt. 3. Heat in roaring Bunsen flame. • Sodium, Na + Yellow • Potassium, K + Lilac • Calcium, Ca 2+ Red • Copper, Cu 2+ Green-blue Precipitation Tests Chloride: add acidified silver nitrate to get a white precipitate if chloride is present. Sulfate: add acidified barium chloride to get a white precipitate if sulfate is present. Carbonate Test 1. Add acid to the sample 2. Pass any gas produced through limewater: will go cloudy if the sample contained carbonate T3: Diamond vs Graphite (HT) Diamond: • Very hard, as all carbon atoms joined with strong covalent bonds. • Used to make cutting tools • Insulator as all electrons locked-tight in bonds, so can’t move. Graphite: • Layers of hexagonal carbon mesh that rub away from each other, as there are only weak forces between the layers. • Used as a lubricant. • Conductor as the electrons between the layers are free to move. This is very rare for a giant covalent structure. T3: Separating Immiscible Liquids •Immiscible = when liquids do not dissolve in each other….like oil and water, one floats on top of the other. •Can be separated with a separating funnel; the denser layer is tapped-off at the bottom. T3: Covalent Bonds • Form when non-metals share electrons between them. • Attraction between each atom and the shared electron pair. • Atoms share electrons to complete their outer shells • One bond is formed for each ‘gap’ in the outer shell • Bonding represented with dot- and-cross diagrams showing only the outer-shell electrons. Example 1: Water Each hydrogen needs one more electron to complete it’s outer shell and the oxygen needs two more. Oxygen forms two single bonds: one to each hydrogen. Example 2: Carbon dioxide (HT only) Carbon needs two more electrons to complete it’s outer shell and each oxygen needs two more. Carbon forms two double bonds: one to each oxygen. H H O O O C T3: Covalent Structures Simple Covalent Molecules • Molecule = A particle made of a small group of atoms, covalently bonded together. • Low melting and boiling point, due to weak attractive forces between molecules.. • Electrical insulator as no electrons free to move. • Examples: water, ammonia, oxygen Giant Covalent •Repeating pattern of many millions of atoms covalently bonded. • High melting/boiling point because much heat energy needed to break strong covalent bonds. • Electrical insulator as no electrons free to move. • Examples: silicon dioxide, diamond, graphite Lead nitrate + potassium iodide lead iodide + potassium nitrate Pb(NO 3 ) 2 (aq) + 2KI(aq) PbI 2 (s) + 2KNO 3 (aq) T3: Separating Miscible Liquids •Miscible = when liquids dissolve in each other…like alcohol and water. •Separate with fractional distillation using a fractionating column. •The components of the mixture have different boiling points, so if you heat it, each component will boil at a different time, allowing you to collect and condense