CHEMISTRY: TOPIC 4 BONDING
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CHEMISTRY: TOPIC 4 BONDING - Marcador
CHEMISTRY: TOPIC 4 BONDING - Detalles
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| What is VSEPR theory? | Pairs of electrons (electron domains) in the valence shell of an atom repel each other so they try to be as far apart as possible to minimise these repulsions |
| What do electron pairs do? | Bonds / lone pairs repel each other and spread apart as far as possible. |
| For the diff bond pairs what is the order of repulsion from most to least | Lone pair - lone pair lone pair - bonding pair bonding - bonding |
| What is electron domain? | Bonds or lone pairs of electrons around an atom in a molecule |
| What counts as on electron domain? | Single bonds double and triple bonds are all 1 DOMAIN not 2, 3 etc |
| What are ions? | A charged particle formed from atoms by a loss or gain of electrons |
| Which groups form cations (positive ions)? | Group 1, 2, 13 = metals |
| Which groups form anions (negative ions)? | Group 15, 16, 17 = non metals |
| What is ionic bonding ? | Forms from an electrostatic force of attraction between oppositely charged ions ( metal - non metal ) |
| What is the structure of an ionic compound and why? (3) | The electrostatic attraction between ions causes them to be surrounded by ions of an opposite charge = ionic lattice fixed arrangement of ions based on a repeating unit usually solid |
| What is coordination number? | Number of ions surrounding a given ion in the lattice eg: in NaCl each Na+ ion is surrounded by Cl- ions and vice cersa = CN is 6 |
| Why don't ionic compounds exist as fixed units? | They have every large numbers of ions and can grow indefinetely so they are expressed as the formula unit which is the simplest ratio |
| What is lattice energy? | Measure of the strength of attraction between the ions within a lattice greater for small and highly charged ions = larger charge density |
| What are physical properties of ionic compounds? (7) | High MP / BP Solid at room temp low volatility usually soluble in ionic or polar solvents but not in non-polar solvents Can conduct electricity in a molten state (liquid) or in an aqueous state (dissolved in water) do not conduct electricity in SOLID state Brittle |
| Why do ionic compounds have high MP / BP? Why are they solid at room temp? | Forces of electrostatic attraction bw oppositely charged ions are very strong and need a lot of heat energy to break throughout the giant lattice the higher the charge the more energy is needed to break the bonds |
| What does ionic compounds MP and BP depend on? | Higher when the charge n the ions is greater because of the increased attraction between the ions |
| What does volatility mean? volatility of ions? | The tendency of a substance to evapourate = ionic compounds have LOW volatility bec the electrostatic forces between the ions are very strong |
| What does solubility mean? | How easily a solute can be disperses through a solvent = solution |
| What does the solubility of ionic compounds depend upon? What are ionic compounds soluble with?(3) | Depends on how much the solute is able to form attractive forces with the solvent = ionic compounds are generally soluble with ionic or polar solvents because they can form bonds with it = soluble in water energy is released when the ions are hydrates (ion-dipole attraction) = energy pays back energy needed to break apart the ionic lattice But not soluble in non polar solvents |
| How do ions get hydrated in water? | Water is polar - seperation of charge in their structure partial charges are attracted to the opposite charges of the ions in the lattice so water molecules dislodge from their position surrounded by water molecules energy is released when ions are hydrated by being surrounded (ion-dipole attractions) |
| Why can ions only conduct electricity when molten or aqueous? (3) | Depends on if the compound can move and carry a charge as a solid the ions are held tightly in the lattice = can't conduct electricity as a liquid state / dissolved in water = ions can move = positive ions moves towards negative electrode and negative ions moves towards the positive electrode = electricity conduction IONS MOVE NOT ELECTRONS |
| Why are ionic compounds brittle? | Movement of ions within the lattice puts ions of the same charge beside each other = repulsive forces causes it to split |
| Trend of gaining and losing electrons? | Tendency to lose elctrons increases DOWN the group 1, 2 13 tendency to gain electrons increases UP the group 17, 16, 15 |
| What is electronegativity? How it affects ionic bonding? | A measure of the atom's ability to attract electrons in a covalent bond bigger electronegativity difference = stronger ionic bond |
| How is a covalent bond held together? (2) | Electrostatic force of attraction between a shared pair of electrons and the positively charged nuclei = non metal - non metal attraction is balanced by the like positive charges of the 2 nuclei = distance |
| What is a molecule? | Group of atoms held together in covalent bonds |
| What is different about transition metals in terms of ionic bonding? | They can form more than one ion |
| What can be used to determine the number of electrons that is lost or gained? | Electron configuration of the atom |
| What are lewis structures? | Lewis structures are diagrams of molecules that represent electron pairs by dots, crosses or lines. They show all of the valence electrons, including lone, or non-bonding pairs. |
| What are lone pairs? | The pairs of electrons in the outer shell that is not involved in covalent bonds |
| How does the number of bonds affect the strength and length of the bond? | Triple bonds is the strongest bonds = strength increases with the number of bonds triple bonds are SHORTER than the other bonds bec of the stronger attraction between the bonding electrons longer the bond is = weaker it is |
| How does the size of the atom affect the bond length? | Electron pair is further away from the nuclei = less strongly attracted = weaker bond |
| Why do noble gases not have electronegativity values? | Because they don't form compounds as they already have a full outer shell |
| How does differences in electronegativities afffect the type of bond formed? | Atoms w similiar electroneg = covalent diff electroneg = ionic |
| What is a coordinate covalent bond? | Both electrons shared come from the same atom which is formed by extra electrons or lone pairs |
| What are resonance structures? | Different structures of the same molecule only differing in the position of the DOUBLE BOND |
| Why do lone pairs have the greatest repulsion? | They are held closer to the central nucleus than bonding pairs so they are therefore closer to the electrons that are bonding than the pairs = stronger repulsion |
| What are polar bonds? | When the electroneg diff bw 2 atoms covalently bonded = electrons are more towards one atom than the other For a molecule to be polar it has to have a positive and negative end |
| How can a molecule be non polar even though the separate bonds are polar? | Although each individual bond is polar due to the difference in electronegativity of the atoms, because of the symmetry of the molecule, the dipoles cancel |
| What are dipole moments? | From the charge and distance bw the charges |
| What are allotropes? (of carbon) (4) | Different forms of the same element diamond, graphite, graphene , fullerene all contain only carbon atoms but join together differently |
| Describe Diamond structure (8) | Giant covalent structure each carbon atom is joined by 4 other carbon atoms = tetrahedral = covalent bonds High MP bec lot of energy needed to break strong covalent bonds very hard bec of strong covalent bonds don't conduct electricity = all electrons held in strong covalent bonds - not free to move Not water soluble bec forces are too strong high thermal condcutivity eg: used as a drill |
| Describe graphite structure: (7) | Giant covalent structure - LAYERED C is covalently bonded to 3 other C = trigonal planar covalent bonds in each layer but only WEAK LONDON FORCES BW LAYERS explains why graphite is a good lubricant = not much force needed to separate layers HIGH MP bec of covalent bonds within layers must be broken NOT water soluble bec of strong covalent bonds between atoms conducts electricity bec of free electrons in 1 carbon atom can move WITHIN layers |
| Describe Graphene: (4) | Single layer of graphite high tensile strength High MP bec of covalent bonds good electrical and thermal conductor bec of C only forming 3 bonds |
| C60 Fullerene (8) | MOLECULAR structure individual C60 molecules with covalent bonds WITHIN molecule LONDON FORCES bw molecules lower MP than graphite and diamond bec ONLY INTERMOLECULAR (LF) have to be overcome when fullerene melts insoluble in water but soluble in some organic solvents (BENZENE) DOES NOT conduct electricity = although there is delocalisation of electrons bec it is MOLECULAR structure = electrons cannot move bw molecules |
| Describe silicon dioxide (SiO2) | Giant covalent structure each silicon atom is bonded to 4 O atoms = tetrahedral High MP bec covalent bonds bw atoms have to be broken eg: clocks and watces |
| How are metals bonded? | Electrostatic attraction between the positive ions in the lattice and the delocalised electrons each electron is attracted by all the positive ions = holds lattice together |
| Describe the structure of a metal | Regular lattice arrangement of positive ions surrounded by a sea of delocalised electrons |
| Reasons why the Mg2+ ion has a higher MP than a Na+ | Mg has a stronger electrostatic attraction bw ions and delocalised e Mg is smaller than Na = delocalised e is closer to the nucleus = more strongly attracted more delocalised electrons per atom in Mg = stronger electrostatic attractions |
| Physical properties of metals? | Lustrous = shiny when cut good electric conductors = Cu used in electrical wires good heat conductors ductile = made into wires malleable = hammered into shape = car bodies |
| Why do metals conduct electricity? | Delocalised electrons that are free to move around |
| Why are metals malleable / ductile? | No direction of their bonding as the metal ions attract the delocalised electrons in all directions = layers can slide over each other exact same as original contrasts with ionic compounds that are brittle as displacement of 1 layer causes like charges to repel |
| What are alloys? | Homogenous mixtures of 2 or more metals or of a metal with a non metal |
| Why are alloys used more commonly? (2) | Stronger and stiffer than pure metals and combine desirable properties of diff metals bec different sized atoms prevents the planes of metal atoms sliding over each other |
| Summary of VSEPR: (4) | Molecule geometry depends on electron pairs in the outer shell of teh central atom 1 multiple bond = 1 electron domain electron pairs (electron domains) in the valence shell of an atom repel each other and will therefore take up positions in space to minimise these repulsions = as far away from each other as possible Lone pairs affect the shape but can't be seen they repel bonding pairs of electrons more than the bonding pairs as they are closer to the bonding electrons |
| What are positive ions called and how are they formed? | Cations are formed by metals losing electrons |
| What are negative ions called? | Anions are formed by non metals gaining electron |
| What does the number of electrons lost or gained depend upon? | The electron configuration of the atom |
| How are ionic compounds held together? | Due to the electrostatic force of attraction between oppositely charged ions |
| What is the electrical conductivity of ions in ionic solutions? | Conduct electrucity because the ions are free to move around |
| What is the solubility of ions in non polar solvents? | Usually not soluble bec a lot of energy is needed to break apart the ionic lattice not paid back by the energy released when the non polar solvent forms interactions with the ions (london forces) |
| What are the polyatomic ions? (10) | NH4+ = ammonia OH- = hydroxide NO3- = nitrate HCO3 - = hydrogen carbonate CO3 (2-) = carbonate SO4(2-) = sulfate PO4(3-) = phosphate O (2-) = oxide S (2-) = sulphide N (3-) = nitride |
| What is the electronegativity difference to be a covalent or ionic molecule? | 0-1.7 = covalent 1.8+ = ionic |
| What are examples of 3 resonance structures? | C6H6 = Benzene CO3 (2-) = carbonate O3 = ozone |
| Why are both bonds in the ozone molecule the same length? (3) | Due to delocalization of pi bond electrons double and single bond = resonance this means that the double bond can change positions which causes resonancy |
| Bond lengths for o-o and o-o double bond? | Found in data booklet (covalent bonds) |
| What does the strength of attraction in covalent bonds depend upon? (2) | Individual atom size nuclear charge |
| What is the octet rule? | How atoms in covalent bonds have a tendency to have a full valence shell with 8 electrons |
| Exceptions to the octet rule? | BF3 = bec boron only has 3 electrons in the outer shell it can only share 3 electrons = max 6 BeCl2 = similiar reason SF6 = has 6 electrons = can form 6 bonds = 12 max (Expanded octet = only p3 and more can expand |
| Explain how NH4+ has a coordinate covalent bond? | NH4+ is formed when a H+ joins with NH3 H+ doesn't have any electrons but N gives a lone pair to it to form a dative bond |
| Explain how H3o+ is formed? | Lone pair is given from the O in H2O to the H+ |
| Example of a lewis acid base reaction? (N and B) | NH3 and BF3 can join together (adduct: 2 molecules joined together) bec boron only has 6 max it can accept a lone pair from N |
| What is the bonding in carbon monoxide and why? | Triple bond bec with a double bond only the Oxygen has 8 electrons carbon still needs 2 more electrons= so oxygen can donate 1 lone pair to carbon = dative bond |
| Explain why there is a 2- charge on CO3 (2-) | Because the carbon can only form 4 bonds so 2 of the oxygen which have single bonds still each need 1 more electron which is given with the negative charge |
| What is the shape, angle and no of domains for BeCl2? | Linear, 180, 2 |
| What is the shape, angle and no of e domains for BF3? | Trigonal planar , 120, 3 |
| What is the shape, angle and no of e domains for CH4? | Tetrahedral, 109.5 , 4 |
| What is the order of repulsion strength for electron pairs? | Lone pair- lone pair > lone pair- bonding pair > bonding pair - bonding pair |
| What is a polar bond? | When the electrons are more towards one atom tan the other but the overall polarity of the molecule depends on the shape polar molecules have an overall dipole moment |
| What makes a molecule non polar/polar? | Made up of same element symmetry = polar noble gases C-H bonds only molecules are non polar H + FON = polar |
| Why are Co2 BF3 and CCl4 non polar even though the bond are polar? | Although each bond is polar due to the diff in electronegativity of the atoms , because of the symmetry of the molecules = dipoles cancel |
| What is the difference between intra and intermolecular bonds? | Intra within the molecule = covalent bonds (Br2) inter between diff molecules = forces that hold it as a liquid but are much weaker than covalent bonds |
| What are the different types of intermolecular bonds? (4) | London dispersion forces = weakest Dipole - dipole interactions dipole-induced dipole interactions hydrogen bonding = strongest |
| What are london dispersion forces? (4) | = much weaker than covalent bonds (in all molecules in solid and liquid states) temporary (instantaneous) dipole - induced dipole interactions get stronger as the number of electrons in molecules increases = relative molecular mass increases explains why it goes from gas to solid down group 17 (forces get stronger as mass increases) = temporary dipoles will be bigger, so more polarisable |
| What are dipole - dipole attraction? (2) | In polar molecules the dipole- dipole and london dispersion forces are stronger than non polar molecules which only has london dispersion forces causes MP and BP to be higher in polar molecules even if they have the same Mr |
| What is hydrogen bonding? | Strongest intermolecular bond but still weaker than a covalent bond When H bonds with very electronegative atoms (F, O , N ) Takes electron density from the hydrogen = partial negative charge while the H would have a partial positive charge molecules with hydrogen bonding are usually soluble in water bec they can H bond to the water (eg ethanol) |
| What is broken when covalent molecular substances are melted? | Only intermolecular bonds |
| What is the general rule for solubility? | Like dissolves like substance will dissolve in solvent if inter forces are similar depends on how much energy is needed to overcome inter forces in solvent and how much energy is released when inter forces are made with teh solvent and solute |