Higher Chemistry Unit 1 Chemical Changes and Structure Summary Notes
Success Criteria I am confident that I understand this and I can apply this to problems? I have some understanding but I need to revise this some more I do not understand this and I need help with it I will be successful if I can Self-Evaluation 1 Describe the importance of controlling reaction rates? x 2 State the two factors required for successful collisions to occur? x 3 State the definition of activation energy? x 4 Name the intermediate structure formed when a successful collision occurs? x 5 6 Describe the effect of increasing or decreasing concentration, pressure, surface area and temperature on reaction rates Use your knowledge of collision theory to explain the effects of concentration, pressure, surface area and temperature on reaction rates? x? x 7 Calculate the relative rate a reaction in s -1? x 8 State the definition of temperature? x 9 Use energy distribution diagrams to explain the effect of changing temperature on the kinetic energy of the particles and reaction rate? x 10 Use potential energy diagrams to calculate the enthalpy change of a reaction? x 11 Use potential energy diagrams to calculate the activation energy of a reaction? x 12 Use potential energy diagrams to determine if a reaction is exothermic or endothermic? x 13 Indicate on a potential energy diagram where the activated complex is formed? x 14 Draw a potential energy diagram? x 15 State the definition of a catalyst? x 16 17 Draw a line on a potential energy diagram to show how the reaction pathway is altered when a catalyst is used Show how the use of a catalyst will effect activation energy on an energy distribution diagram? x? x 18 State the definition of covalent radii? x 19 Identify the trends in covalent radii across a period and down a group in the periodic table? x
20 Explain the trends in covalent radii across a period and down a group in the periodic table? x 21 State the definition of electronegativity? x 22 23 Identify the trends in electronegativity across a period and down a group in the periodic table Explain the trends in electronegativity across a period and down a group in the periodic table? x? x 24 State the definition of ionisation energy? x 25 Write ionisation energy equations? x 26 27 Identify the trends in ionisation energy across a period and down a group in the periodic table Explain the trends in ionisation energy across a period and down a group in the periodic table? x? x 28 Name the three types of intramolecular bonds? x 29 Explain, in terms of electronegativity, how a pure covalent (non-polar) bond is formed? x 30 Explain, in terms of electronegativity, how a polar covalent bond is formed? x 31 Identify which element in a polar covalent bond will be partially negative and which will be partially positive? x 32 Indicate partial charges within a polar covalent bond? x 33 Describe how the difference in electronegativity will affect the degree of polarity of a bond? x 34 Explain how an ionic bond is formed? x 35 Identify which element in an ionic bond will form a negative ion and which will form a positive ion? x 36 Indicate charges within an ionic bond? x 37 Indicate the position of different types of bonds within the bonding continuum? x 38 Describe the properties of covalent and ionic bonding? x 39 Name the three types of intermolecular bonds (van der Waals forces)? x 40 State the definition of London dispersion forces? x 41 Explain how permanent dipole- permanent dipole interactions are formed? x 42 Explain how hydrogen bonds are formed? x 43 Describe how the presence of different intermolecular bonding affects the properties of the substance e.g. melting point, boiling point and viscosity? x
44 Compare the strength of different types of intermolecular bonds? x 45 Describe how the type of bonding present (ionic, polar or non-polar) can affect the solubility of a substance? x
Key Area 1.1 Controlling the Rate It is important when carrying out industrial processes to understand how to control the rate of a chemical reaction o This can increase the volume/mass etc. of product produced o This can reduce the safety risks e.g. the probability of explosion For a reaction to be successful, the particles of the reactants must collide o They must overcome the activation energy (EA) o They must collide with the correct collision geometry When reactant particles collide with the activation energy (EA) and correct collision geometry an activated complex is formed. o An activated complex is an unstable intermediate that can break down to form the products of the reaction. Variables Affecting Reaction Rate There are many variables that can affect the rate of a reaction Changing concentration, temperature, pressure and surface area can affect the rate of reaction The following factors can be changed to increase the rate of a reaction 1. Increasing the concentration of a reactant increases the number of particles in the same volume of space, increasing the probability of successful collisions 2. Increasing temperature results in an increase in energy, more particles can overcome activation energy 3. Increasing the pressure of the reaction decreases the space the particles have to move, increasing the probability of collisions 4. Increasing the surface area (decreasing particle size), increases the probability of successful collisions The opposite of each is true when decreasing the rate of reaction Calculating the Rate The relative rate of a reaction (s -1 ) can be calculated as follows:
Energy Distribution Diagrams Temperature is a measure of the average kinetic energy or speed of the particles of a substance At any given temperature, the particles of a substance will have a range of kinetic energies o This can be shown using an energy distribution diagram These diagrams can give us information regarding the number of particles that overcome activation energy and the average kinetic energy of the particles Energy distribution diagrams can be affected by temperature o Increasing the temperature results in particles having a greater kinetic energies and more will have energies greater than the activation energy o It is important to understand how this would affect the distribution curve Potential Energy Diagrams The energy change between reactants and products can be shown using a potential energy diagram Important information can be obtained from a potential energy diagram o Activation energy of the forward reaction o Activation energy of the reverse reactions o Change in enthalpy o If the reaction is exothermic or endothermic
The diagram below is an example of a potential energy diagram o There is an energy gain from reactants to products o An energy gain means this reaction is endothermic o The change in enthalpy will be positive (+ H) Activation energy for forward reaction Activation energy for reverse reaction Enthalpy change (+ H) The diagram below is an example of a potential energy diagram o There is an energy loss from reactants to products o An energy loss means this reaction is exothermic o The change in enthalpy will be negative (- H) Activation energy for forward reaction Activation energy for reverse reaction Enthalpy change (- H) The enthalpy change ( H) of a reaction is the difference in potential energy between reactants and products o The units of enthalpy change ( H) are kjmoll -1 o H = H (products) H (reactants) o Exothermic reactions have a negative H, endothermic reactions have a positive H
Catalysts Catalysts supply an alternative reaction pathway with a lower activation energy (EA) o This results in a change to the potential energy diagram Activation energy for the uncatalysed forward reaction Activation energy for the catalysed forward reaction Key Area 1.2 Periodicity There are three main trends on the periodic table o Covalent Radii o Electronegativity o Ionisation energy Covalent Radii The covalent radius of an atom is a measure of the size of the atom o Covalent radii decrease across a period o Covalent radii increase down a group The radius of an atom decreases across a period due an increase in nuclear charge o Nuclear charge is directly related to the total number of protons in an atom Nuclear charge increases as the number of protons increases Nuclear charge decreases as the number of protons decreases o As nuclear charge increases, the electrons are pulled more closely to the nucleus resulting in a decrease in radius The radius of an atom increases down a group due to an increase in the number of electron shells and shielding o Shielding occurs when the outer electrons are blocked from the nucleus by the inner electrons resulting in the outer electrons being less strongly attracted to the nucleus Shielding increases as the number of filled electron shells increases Shielding decreases as the number of filled electron shells decreases o As the number of shells increases, shielding occurs resulting in an increase in radius
Electronegativity Electronegativity is the ability of an atom to attract electrons o High electronegativity, strong pull on electrons o Low electronegativity, weak pull on electrons Electronegativity increases across a period Electronegativity decreases down a group Electronegativity increases across a period due an increase in nuclear charge o As the nuclear charge increases the nucleus can pull electrons more strongly, resulting in an increase in electronegativity Electronegativity decreases down a group due to an increase in the number of electron shells and shielding o As the number of electrons increases, more shells become full and shielding occurs, resulting in a decrease in electronegativity Ionisation Energy Ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state o First ionisation energy is the energy required to remove the first electron o Second ionisation energy is the energy required to remove a second electron Ionisation equations can be written using the following formula o First ionisation energy: E(g) E + (g) + e - o Second ionisation energy: E + (g) E 2+ (g) + e - o For example, the equation or the fist ionisation of sodium is Na(g) Na + (g) + e - Ionisation energy increases across a period due to an increase in nuclear charge o As the nuclear charge increases the electrons are pulled more closely to the nucleus, more energy is required to remove the electron, resulting in an increase in ionisation energy Ionisation energy decrease down a group due to an increase in the number of electron shells and shielding o As the number of electrons increases, more shells become full and shielding occurs. The outer electrons are held less tightly to the nucleus and require less energy to remove, lowering the ionisation energy. When explaining why. o An increase in nuclear charge should always explain the trend across a period o An increase in the number of shells and shielding should always explain the trend down a group
Key Area 1.3 Structure and Bonding There are two main categories of bonding o Intramolecular bonding The bonding that occurs between atoms o Intermolecular bonding The bonding that occurs between molecules Intramolecular bonding There are three types of intramolecular bonds o Pure covalent (non-polar) o Polar covalent o Ionic Pure covalent (non-polar) A pure covalent bond is formed when two atoms share a pair of electrons equally o both atoms have the same electronegativity o diatomic elements have pure covalent bonds Electronegativity 2.2 Electronegativity 2.2 No difference in electronegativity so a pure covalent bond is formed Polar covalent A polar covalent bond is formed when two atoms share a pair of electrons unequally o The atoms have a difference in electronegativity The atom with the highest electronegativity becomes partially negative () The atom with the lowest electronegativity becomes partially positive (δ+) o For example, the bond between carbon and chlorine Carbon has an electronegativity of 2.5, this is less than chlorine result in carbon becoming partially positive δ+ Chlorine has an electronegativity of 3.0, this is greater than carbon result in chlorine becoming partially negative Exception to the rule o Carbon to hydrogen bonds are always considered to be non-polar
Ionic An ionic bond is formed when electrons are transferred from one atom to another resulting in the formation of ions o The atom with the lowest electronegativity forms a positive ion o The atom with the highest electronegativity forms a negative ion o The bond is formed due to the electrostatic attraction between the positive and negative ions o For example, the bond between sodium and chlorine in sodium chloride Sodium has an electronegativity of 0.8, this is less than chlorine resulting in sodium becoming a positive ion + - Chlorine has an electronegativity of 3.0, this is greater than sodium resulting in chlorine becoming a negative ion When a metal atom is bonded to a non-metal atom, these are most commonly ionic bonds o Metal atoms form positive ions when they lose electrons as they have lower electronegativities o Non-metals form negative ions when they gain electrons as they have higher electornegativities The Bonding Continuum Pure covalent bonding and ionic bonding can be considered as being at opposite ends of a bonding continuum Pure Covalent Polar Covalent Ionic Intermolecular bonding Intermolecular bonds are known as van der Waals forces There are three types of van der Waals forces o London dispersion forces o Permanent dipole- permanent dipole interactions o Hydrogen bonding London dispersion forces London dispersion forces are forces of attraction that can occur between all atoms and molecules o London dispersion forces occur due to the presence of temporary/induced dipoles that form when electrons move in atoms and molecules o London dispersion forces are the weakest of all van der Waals forces
Permanent dipole-permanent dipole interactions Permanent dipole-permanent dipole interactions occur between polar molecules o Polar molecules are molecules that contain polar bonds o The interaction occurs when a partially negative charge within one molecule is attracted to the partially positive charge of another molecule o For example, trichloromethane (CHCl3) Carbon has an electronegativity of 2.5 and chlorine has an electronegativity of 3 There is a different in electronegativity resulting in a polar bond being formed δ+ Carbon-hydrogen bonds are considered to be non-polar Chlorine has the higher electronegativity and has a partial negative charge Carbon has the lower electronegativity and has a partial positive charge The partially negative chlorine in one molecule of trichloromethane is electrostatically attracted to the partially positive carbon in another molecule, forming permanent dipolepermanent dipole interactions Permanent dipole- permanent dipole interactions are stronger than London dispersion forces
Hydrogen Bonding Hydrogen bonds can occur between molecules if they contain hydrogen bonded to highly electronegative atoms e.g. nitrogen, oxygen or fluorine o When a hydrogen atom is bonded to a highly electronegative atom is becomes partially positive and the other atom becomes partially negative o The positive hydrogen atom of one atom is electrostatically attracted to the partial negative atom in another molecule Oxygen has an electronegativity of 3 and hydrogen has an electronegativity of 2.2 There is a different in electronegativity resulting in a polar bond being formed Oxygen has the higher electronegativity and has a partial negative charge Hydrogen has the lower electronegativity and has a partial positive charge The partially negative oxygen in one molecule of water is electrostatically attracted to the partially positive hydrogen in another molecule, forming hydrogen bonds Hydrogen bonds are the strongest of all the van der Waals forces Strength of intermolecular bonds
Properties The type of intermolecular bonds present can affect the properties of a substance o Melting and boiling point o Viscosity o Solubility Melting and boiling point The melting and boiling point of a substance is dependent on the energy required to break the intermolecular bonds present o The stronger the intermolecular bonds, the more energy required, resulting in higher melting and boiling points If a substance contains hydrogen bonding, the melting and boiling point of that substance will be greater than a substance with only weak London dispersion forces between molecules Viscosity The strength and number of intermolecular forces between molecules can affect the viscosity of a substance o The stronger the intermolecular forces, the more viscous the substance will be o The greater number of intermolecular forces, the more viscous the substance will be Carbon to hydrogen bonds are non-polar The only intermolecular forces present are weak London dispersion forces Pentane is less viscous than ethanol Ethanol contains polar bonds Hydrogen bonding can form between the molecules δ+ δ+ Ethanol is more viscous than hexane δ+ 1,3-ethan-diol contains polar bonds Hydrogen bonding can form between the molecules 1,3-ethan-diol is more viscous than ethanol as it has more hydrogen bonding δ+ δ+ δ+
Solubility The solubility of a substance is dependent on the polarity of the substance and the polarity of the solvent o Polar substances will dissolve in polar solvents o Non-polar substances will dissolve in non-polar solvents o Like Dissolves Like Special Examples Some molecules contain polar bonds, however, the molecule is non-polar overall o The polar bonds cancel each other out o Examples include carbon dioxide, carbon tetrachloride and carbon tetrafluoride δ+ δ+ δ+ Graphite and fullerene are structures that only contain carbon atoms o The structure and bonding of these result in graphite having a higher melting point than fullerene The carbon atoms in fullerene form discrete molecule structures containing, for example, 60 carbon atoms (C60) o There are strong non-polar, covalent bonds between these carbon atoms o As there are no polar bonds, the intermolecular forces present between the molecules are London dispersion forces o London dispersion forces are weak and do not require high amounts of energy to break This results in fullerene having a low melting point The carbon atoms in graphite form layers of carbon atoms bonded to form a covalent network structure o There are strong non-polar, covalent bonds throughout the structure o There are delocalised electrons between the layers of carbon atoms o To melt this structure you would need to break the layers of strongly bonded carbon atoms which requires high amounts of energy This results in fullerene having a high melting point