1. Reactions can be followed by measuring changes in concentration, mass and volume of reactants and products.

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1 Higher Chemistry - Traffic Lights Unit 1 CHEMICAL CHANGES AND STRUCTURE I know: Controlling the rate Collision theory and relative rates 1. Reactions can be followed by measuring changes in concentration, mass and volume of reactants and products. 2. The average rate of a reaction, or stage in a reaction, can be calculated from initial and final quantities and the time interval. 3. The rate of a reaction, or stage in a reaction, is proportional to the reciprocal of the time taken (1/t). 4. Reaction rates can be controlled by chemists. 5. Very slow reactions are not economically viable in manufacturing processes. 6. Reaction rates that are too fast risk causing a thermal explosion. 7. Collision theory can explain the effects of concentration, pressure, surface area (particle size), temperature and collision geometry on reaction rates. 8. Temperature is a measure of the average kinetic energy of the particles of a substance. 9. The activation energy is the minimum kinetic energy required by colliding particles before reaction will occur. 10. Energy distribution diagrams can be used to explain the effect of changing temperature on the kinetic energy of particles. 11. The effect of temperature on reaction rate can be explained in terms of an increase in the number of particles with energy greater than the activation energy.

2 12. With some chemical reactions, light can be used to increase the number of particles with energy greater than the activation energy. Potential energy diagrams Reaction pathways and enthalpy changes 1. A potential energy diagram can be used to show the energy pathway for a reaction. 2. The enthalpy change is the energy difference between products and reactants. 3. The enthalpy change can be calculated from a potential energy diagram. 4. Exothermic changes cause heat to be released to the surroundings; endothermic changes cause absorption of heat from the surroundings. 5. The enthalpy change has a negative value for exothermic reactions and a positive value for endothermic reactions. 6. The activated complex is an unstable arrangement of atoms formed at the maximum of the potential energy barrier during a reaction. 7. The activation energy is the energy required by colliding particles to form an activated complex. 8. The activation energy can be calculated from potential energy diagrams. Catalysts 1. A catalyst provides an alternative pathway with a lower activation energy. 2. A potential energy diagram can be used to show the effect of a catalyst on activation energy. 3. Catalysts can be classified as either homogeneous or heterogeneous. 4. Heterogeneous catalysts work by the adsorption of reactant particles. 5. The surface activity of a catalyst can be reduced by poisoning. 6. Impurities in the reactants result in the industrial catalysts having to be regenerated or renewed. 7. Catalytic converters are fitted in cars to catalyse the conversion of poisonous carbon monoxide and oxides of nitrogen to carbon dioxide and nitrogen.

3 8. Cars with catalytic converters only use 'lead-free' petrol to prevent poisoning of the catalyst. 9. Enzymes catalyse the chemical reactions which take place in the living cells of plants and animals and are used in many industrial processes. Periodicity Elements and Trends 1. The modern Periodic Table is based on the work of Mendeleev who arranged the known elements in order of increasing atomic masses in conjunction with similar chemical properties, leaving gaps for yet to be discovered elements. 2. There are variations in the densities, melting points and boiling points of the elements across a period and down a group. 3. The covalent radius is a measure of the size of an atom; it decreases across a period and increases down a group. 4. The trends in covalent radius across periods and down groups can be explained in terms of the number of occupied shells and the nuclear charge. 5. Atoms of different elements have different attractions for bonding electrons. 6. Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond. 7. Electronegativity values increase across a period and decrease down a group. 8. Electronegativity trends can be rationalized in terms of nuclear charge, covalent radius and the presence of screening inner electrons. 9. The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. 10. The second and subsequent ionisation energies refer to the energies required to remove further moles of electrons. 11. The trends in the first ionisation energies across periods and down groups can be explained in terms of the atomic size, nuclear charge and the screening effect due to inner shell electrons.

4 Structure and bonding First 20 elements and bonding types 1. The first 20 elements of the Periodic Table can be categorised according to bonding and structure: metallic (Li, Be, Na, Mg, Al, K, Ca) covalent molecular solids (P4, S8, and C (fullerenes)) covalent molecular gases (H2, N2, O2, F2, C12,) covalent network (B, C (diamond, graphite), Si) monatomic (noble gases) 2. A metallic structure consists of a giant lattice of positively charged ions and delocalised electrons. 3. Metallic bonding is the electrostatic force of attraction between positively charged ions and delocalised outer electrons. 4. In a covalent bond, atoms share pairs of electrons. 5. The covalent bond is a result of two positive nuclei being held together by their common attraction for the shared pair of electrons. 6. A covalent molecular structure consists of discrete molecules held together by weak intermolecular forces. 7. A covalent network structure consists of a giant lattice of covalently bonded atoms. 8. A monatomic structure consists of discrete atoms held together by van der Waals' forces. 9. Polar covalent bonds are formed when the attraction of the atoms for the pair of bonding electrons is different. 10. Delta ( + ) positive and delta negative ( - ) notation can be used to indicate the partial charges on atoms, which give rise to a dipole. 11. Pure covalent bonding and ionic bonding can be considered as being at opposite ends of a bonding continuum with polar covalent bonding lying between these two extremes. 12. The larger the difference in electronegativities between bonded atoms, the more polar the bond will be.

5 13. If the electronegativity difference is large then the movement of bonding electrons from the element of lower electronegativity to the element of higher electronegativity is complete, resulting in the formation of ions. 14. Compounds formed between metals and non-metals are often, but not always, ionic. 15. An ionic structure consists of a giant lattice of oppositely charged ions. 16. Ionic bonding is the electrostatic attraction between positively and negatively charged ions. 17. Compounds each adopt one of three structures in the solid state: covalent molecular covalent network, including silicon dioxide and silicon carbide ionic 18. The type of bonding in a compound is related to the position of its constituent elements in the Periodic Table. Intermolecular forces of attraction 1. All molecular elements and compounds and monatomic elements condense and freeze at sufficiently low temperatures. 2. For this to occur, some attractive forces must exist between the molecules or discrete atoms. 3. Any intermolecular forces acting between molecules are known as van der Waals forces. 4. There are several different types of van der Waals forces such as London dispersion forces and permanent dipole: permanent dipole interactions, which include hydrogen bonding. 5. London dispersion forces are forces of attraction that can operate between all atoms and molecules. 6. London dispersion forces are much weaker than all other types of bonding. 7. London dispersion forces are formed as a result of electrostatic attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules.

6 8. London dispersion forces are increase in strength as the number of electrons increases within an atom or molecule. 9. A molecule is described as polar if it has a permanent dipole. 10. The spatial arrangement of polar covalent bonds can result in a molecule being polar. 11. Permanent dipole-permanent dipole interactions are additional electrostatic forces of attraction between polar molecules. 12. Permanent dipole-permanent dipole interactions are stronger than London dispersion forces for molecules of equivalent size. 13. Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegative element such as fluorine, oxygen or nitrogen, are highly polar. 14. Hydrogen bonds are electrostatic forces of attraction between molecules which contain these highly polar bonds. 15. A hydrogen bond is stronger than other forms of permanent dipole-permanent dipole interaction but weaker than a covalent bond. Properties 1. Melting points, boiling points and viscosity can all be rationalized in terms of the nature and strength of the intermolecular forces which exist between molecules. 2. By considering the polarity and number of electrons present in molecules, it is possible to make qualitative predictions of the strength of the intermolecular forces. 3. The melting and boiling points of polar substances are higher than the melting and boiling points of non-polar substances with similar numbers of electrons. 4. The anomalous boiling points of ammonia, water and hydrogen fluoride are a result of hydrogen bonding. 5. Boiling points, melting points, viscosity and miscibility in water are properties of substances which are affected by hydrogen bonding. 6. Hydrogen bonding between molecules in ice results in an expanded structure which causes the density of ice to be less than that of water at low temperatures.

7 7. Ionic compounds and polar molecular compounds tend to be soluble in polar solvents such as water and insoluble in non-polar solvents. 8. Non-polar molecular substances tend to be soluble in non-polar solvents and insoluble in polar solvents. 9. The uses of diamond, graphite and silicon carbide are related to their structures and properties. 10. Fullerenes are the subject of current research and applications are being sought.