Chemistry Outline Yr10

CHEMISTRY Focus 1. Matter is compound of atoms, which combine to make substances 1.1 Recall definition of an atom, element, mixture, and molecule and use diagrams to enhance the definitions • Atom: Smallest unit of matter that can exist by itself. Each atom has a central nucleus with electrons orbiting within electron shells. The nucleus is made of protons and neutrons and gives the atom most of its mass. • Element: Pure substance that cannot be broken down chemically and consists of one atom type only • Compound: Pure substance that consists of two or more elements that are chemically joined in fixed rations to each other • Mixture: Impure substance consisting of two or more elements or compounds that are not chemically united and do not exist in fixed proportions to each other • Molecule: Smallest unit of a substance that is made up of one or more atoms 1.2 Describe how an understanding of the atom has evolved over time to the present day model Name Year Discovery / Proposal Democritus (470Put forward the idea that matter could be divided into smaller particles until it 360BC) became small, hard, indivisible particles, which he called atoms John Dalton 1803 Proposed that matter is made of indivisible particles or atoms and that each element is made up of its own characteristic atoms. Compounds contain atoms of different elements joined in a simple ratio (‘Ball’ model) Sir John 1904 Discovered the electrons and showed that electrons are negatively charged Thomson particles and smaller and lighter that the smallest atom, hydrogen. Proposed that the atom has electrons embedded in a nucleus of positive charge (‘Plum Pudding’ model) Sir Ernest 1911 Identifies positively charged particles (protons), which gave the atom most of its Rutherford mass. He proposed that an atom had a small positive nucleus that was surrounded by and equal number of orbiting electrons (‘Empty Space’ model) Niels Bohr 1913 Proposed an atomic model where electrons travel around the nucleus in fixed orbits and each orbit has a definite amount of energy (‘Orbital’ model) Sir James 1932 Discovered the neutron Chadwick Found it to be of similar mass to the proton Erwin 1962 Undertook research that indicated that electrons travel randomly in a region or Schrödinger space round the nucleus rather than in fixed orbits. Each space (electron shell) & Werner has a definite energy level. Energy levels have sublevels, or orbitals, within Heisenberg them (‘Electron Cloud’ model) Focus 2. Atoms are made of smaller particles and have a common structure 2.1 Construct a table to indicate the features of the main sub-atomic particles mass, charge, and location relative to nucleus Particle in atom Charge of particle Location of particle Proton Positive Nucleus Neutron Neutral Nucleus Electron Negative Orbits nucleus in shell

including relative size, Mass of Particle One unit One unit 1/1836

2.2 Identify that electrons orbit the nucleus in energy levels or shells and draw diagrams of the first 20 elements to indicate the number of electrons in each shell Electrons exist in ‘shells’ or ‘energy levels’ around the nucleus Shell #1 – 2 electrons Shell #3 – 8 electrons Shell #2 – 8 electrons Shell #4 – 18 electrons Full shelled atoms: not reactive

Focus 3. Atoms differ from each other in terms of their subatomic particles 3.1 Defer atomic number, atomic weight and valency, and locate this information from the periodic table • Atomic number: the number of protons in an atom’s nucleus • Atomic weight: mass of one mole of an element • Mass number: the number of protons plus the number of neutrons. Different isotopes of an element have different numbers of neutrons, so they have different mass numbers. • Valency: charge of an atom when It becomes an ion 3.2 Determine the number of protons, neutrons and electrons using the periodic table • Atomic number = protons in nucleus = number of electrons • Atomic mass = number of protons + number of neutrons 3.3 Describe an isotope as having the same atomic number but differing numbers of neutrons • Isotope: Version of an atom, which, in its nuclei, has the same amount of protons and different amount of neutrons, resulting in different masses 3.4 Define valency and relate this to electron shells • Valency: the number of electrons that an atom needs to gain, lose or share in a chemical reaction to complete its electron shell. It is sometimes called the ‘combining power’ of an atom. Some elements may have more than one valency, because they combine in different ways Focus 4. The periodic table is used to classify elements and display relationships between them 4.1 Describe the following relationships on the periodic table: groups, periods, metals, non metals and metalloids, reactivity and radioactivity • Reactivity: Different metals have different abilities to react with other substances. The reactivity of a metal can be determined by placing them in cold/hot acid or using displacement reactions. When a metal is placed in an aqueous solution that contains ions from a less active metal, the less active metal will come out of the solution and the more active metal will go into the solution. • Radioactivity: the breakdown or decay of unstable substances. They release radiation in the form of alpha particles or beta particles and/or nuclear energy in the form of gamma rays. Radioactive elements include elements 58-71 and 90-103 4.2 Describe the relationship between electron arrangement and the organization of the periodic table • Groups: vertical columns on the periodic table o All elements in a group behave similarly in chemical reactions and have the same valency. Number of shells increases but valency remains the same • Periods: horizontal rows on the period table. o The numbers of shells remains the same but the number of electrons increases across the period. Focus 5. Atoms for compounds by joining together and the outer electrons govern how this occurs 5.1 State the law of conservation of matter • Matter cannot be created, lost or destroyed, during physical or chemical changes; it can only be transformed from one form to another 5.2 Describe how a compound can be formed • A compound is a pure substance made up of two or more different elements in a fixed ratio

5.3 Explain compound formation in terms of Loss and gain of electrons: • Ionic compounds

Sharing of electrons: • Covalent compounds

5.4 Describe how ionic and covalent compounds are formed using example • Covalent bonds: o Formed when the atoms share electrons o By sharing electrons each atom obtains a full outer electron shell, subsequently becoming stable o Electron dot diagrams often used to represent covalent compounds o Covalent compounds contain individual molecules • Ionic bonds: o Formed when one or more electrons is transferred from one atom to another o Each atom receives or loses electrons and in this way obtains a full outer electron shell and becomes stable o In the process the atoms become charged and form ions o An attractive force or bond occurs between ions with opposite charge and holds the ions together. As a result, they form an ionic compound. • e.g. Na+ & Clo Sodium atom has 1 electron in its outer shell and is unstable o Chlorine atom has 7 electrons in its outer shell but needs 8 to become stable • Sodium gives 1 electron to chlorine so that both atoms become ions and now have full outer shells and stable formation Focus 6. A wide range of substances exist and they can be classified based on common properties 6.1 Define the term polyatomic ion • Polyatomic ion: A group of atoms with a charge (radicals)

6.2 Identify common elements, salts, acids and bases by their common names, systematic names and chemical formulae Radical Formula Ammonium NH4+ Hydroxide OHCarbonate CO32Nitrate NO3Sulfate SO42Sulfite SO32Phosphate PO43-

Results from addition of Carbonic acid (H2CO3) Nitric acid (HNO3) Sulfuric acid (H2SO4) Sulfurous (H2SO3) Phosphoric acid (H3PO4)

Focus 7. Chemical substances are names systematically based on the elements they contain. 7.1 Apply rules to name simple chemical compounds • Metal - non-metal compound; o Two elements only: • Metallic element named first • Name of non-metal is shortened • Suffix ‘-ide’ is added to this shortened name o Three or more elements: • Metallic element is named first • Chemical radical named second • Non-metal – non-metal compound; o If hydrogen is present it is named first

If no hydrogen is present, solid non-metal is named first Only two elements: • Name of second is shortened and suffix ‘-ide’ is added to the shortened form • Prefixes ‘mon-’, ‘di-’, ‘tri-’, ‘tetra-’ etc. are used on the name of the second to indicate how many atoms of it there are in the formula o More than two elements: • First is named and is followed by the name of the radical 7.2 Identify that compounds are classified into groups based on chemical characteristics • pH <7……acid(red UI) • pH 7 ……neutral(greenUI) • pH >7……base(blue UI) o o

7.3 Describe the characteristics of compound Compound Acid Must Contain H+ ions Character Sour when dilute -istics Corrosive when concentrated Hydrogen ion is the donator Common HCl

groups including acids, bases and salts Salt Base Metal + non metal OH- or NH3 (ammonia) Distinct crystal shapes Bitter and soapy when dilute Soluble in water, Corrosive forming electrolyte Hydrogen ion is the acceptor solution salts=some radicals/polyatomic NaCl OH-

7.4 Identify the structure of an evaluation and write a short evaluation Focus 8. Different types of reactions exist and these have general forms and specific forms that can be represented by equations 8.1 Define corrosion, neutralisation • Corrosion: deterioration of a material that results from a reaction with its environment. Occurs when a metal reacts with oxygen in the presence of water and becomes ionised • Neutralisation: Occurs when an acidic solution is mixed with an alkaline (base) solution and the number of H+ ions balances the number of OH- ions

8.2 Identify the general form of reactions involving acids on metals and carbonates Acidic reactions occur with metals and ionic compounds (not covalent compounds) • Acid + Metal  Salt + Hydrogen gas • Acid + Carbonate  Salt + Water + Carbon dioxide gas • Acid + Base  Salt + Water (neutralisation) 8.3 Write general word equations for each reaction Decomposition • The breakdown of a compound into simpler substances o e.g. CuCO3  CuO + CO2 Combustion • The burning of any substance in air to produce energy o e.g. 2Mg + O2  2MgO + light energy Corrosion • The reaction of a metal with oxygen in the presence of water. This eats away the metal o e.g. 4Fe + 3O2 (water) 2Fe2O3 Precipitation • Reactions that produce a solid from the mixing of two solutions o e.g. AB + CD  BC + AD 8.4 Write word equations for specific examples of each of these reactions Decomposition

Copper Carbonate (heat) Copper oxide + Carbon dioxide Combustion Magnesium + Oxygen  Magnesium oxide + light energy Corrosion Iron + Oxygen + Water  Hydrated Iron Oxide (rust) Precipitation Lead iodide + potassium nitrate  Potassium iodide + lead nitrate 8.5 Write balanced chemical equations for specific exampled of each of these reactions Focus 9. Reactions can be represented with balanced chemical equations which provide very specific information about the chemical involved 9.1 Construct word equations from observations and written descriptions of a range of chemical reactions Focus 10. Different elements have different abilities to react, and this governs how they are used 10.1 Define chemical activity • Chemical activity: A measure of how different elements/compounds in a non-ideal gas or solution, interact with each other 10.2 Extract information from data tables to sort metals in order of activity Metal Reaction With Water Reaction With Dilute Acid Potassium Violent with cold water Violent Sodium Calcium Moderate with cold water Magnesium Burns in steam Moderate Aluminium Protected by oxide film Zinc

Fast with steam

Iron

Reversible

Tin Lead (Hydrogen) Copper

Slow No Reaction

Mercury No Reaction Silver Gold Platinum NB. • When two metals are in contact, the more reactive metal corrodes more readily compared to the less reactive metal. It is thus being sacrificed. • When a metal is placed in an aqueous solution that contains metal ions from a less active metal, the less active metal will come out of the solution and the more active metal will go into the solution. This is called displacement. • You can tell this has happened because coating forms on the metal placed into the solution Examples of Displacement An iron nail is dipped into a copper sulphate solution, a copper coating forms on the iron nail An aluminium wire is dipped into a zinc sulphate solution, a zinc coating forms of the aluminium wire 10.3 Relate activity of metals to their use Sacrificial Anodes

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In Plumbing – Joining pieces of iron or tin to copper or brass fittings that will be connected to an iron pipe • To protect iron pipes and steel wharves – A block of magnesium is connected to the iron or steel, and it will corrode but the iron it is protecting will not. • In Boating – A piece of magnesium or zinc is attached to an outboard motor to protect its steel parts. The magnesium has to be replaced eventually, because corrosion occurs Galvanising • Coating the iron with zinc, a more reactive metal. Even though the zinc corrodes, it protects the iron • “painted on” sacrificial anode