bonding [P1] PAPER 1+2 Flashcards
why do lone pairs of electrons influence the shape/bond angles in a molecule?
• electron pairs repel to be as far as possible
• lone pairs repel more than bond pairs
define the term ‘electronegativity’
power of an atom to attract the pair of electrons in a covalent bond
how do permanent dipole-dipole forces between molecules arise?
• difference in electronegativity leads to bond polarity
• the dipoles don’t cancel each other out, so the molecule has an overall permanent dipole
• there is an attraction between the δ+ on one molecule and the δ- on another
how do van der Waals’ forces between molecules arise?
• electron movement in the first molecule (OR temporary dipole in the first molecule) induces a dipole in another molecule
• induced-temporary attraction OR δ+ attracts δ- in adjacent molecules
put the following intermolecular forces in order of strength, from most to least strongest: dipole-dipole, van der Waal’s, hydrogen bonding
• hydrogen bonding
• dipole-dipole forces
• van der Waal’s forces
⚠️ ONLY simple covalent molecules have IMF forces
describe the bonding in metals
giant metallic lattice held together by strong electrostatic forces of attraction between positive metal ions and delocalised electrons (metallic bonds)
why do metals have a high melting point?
• giant metallic lattice held together by strong electrostatic forces of attraction between positive metal ions and delocalised electrons
• lots of energy required to overcome these forces
why are metals ductile and malleable?
layers of ions are able to slide over each other (while still being held together by delocalised electrons)
draw a diagram to show the arrangement of particles in a crystal of a metal
• regular pattern of identical positive ions (draw minimum of 6)
• delocalised electrons (same number of them as the overall positive charge, e.g. if the ion is 2+ then two delocalised electrons per ion)
why are metals good conductors of heat?
the tightly packed ions efficiently pass vibrational energy on effectively through the lattice structure
why are metals good conductors of electricity?
free/mobile delocalised electrons are able to move through the lattice structure
what does the strength of a metallic bond depend on?
• charge of ion
• size of ion
explain why magnesium has a higher melting point than sodium
[Mg and Na are in different groups, so…]
• same shielding
• magnesium ions have a higher charge than sodium ions
• so stronger forces of attraction between delocalised electrons and positive Mg ion OR metallic bonding is stronger in Mg
explain why strontium has a higher melting point than barium
[they are in the same group, so…]
• both ions have the same charge
• strontium ions are smaller ions
• strontium ions have less shielding
• stronger electrostatic forces of attraction between the positive ions and delocalised electrons in a strontium ion, so stronger metallic bonding (so more energy needed to overcome these forces)
describe the bonding in an ionic compound
giant ionic lattice held together by many strong electrostatic forces of attraction between oppositely charged ions
why do ionic compounds have high melting points?
• giant ionic lattice held together by many strong electrostatic forces of attraction between oppositely charged ions
• lots of energy required to overcome these forces
in what state can ionic compounds conduct electricity and why?
• molten or dissolved
• ions are free to move and carry charge
deduce why the type of bonding in nitrogen oxide is covalent rather than ionic
small difference in electronegativity
suggest why sodium iodide has a lower melting point than sodium bromide
iodide is a larger ion (so weaker attraction to the sodium ion)
common polyatomic ions - ammonium, cyanide, hydroxide, sulfate, nitrate, carbonate, phosphate, chlorate
what is a covalent bond?
shared pairs of electrons
what is a dative covalent bond?
shared pair of electrons with both electrons supplied by one atom
name the type of bond formed when a molecule of BF4- reacts with an F- ion and explain how
• dative covalent
• lone pair of electrons on F- are donated to BF4-
how is a dative covalent bond represented on a diagram?
arrow from atom donating electron pair to other atom
e.g. lone pair of electrons is donated from N of NH3 to H+ to form NH4+ :
how do hydrogen bonds arise?
attraction between lone pairs of electrons on N/O/F (that is bonded to H) and H δ+ (bonded to N/O/F) of the adjacent molecule
what must be present for hydrogen bonding to occur?
hydrogen bonded with nitrogen, oxygen or fluorine in the molecule
hydrogen bonding diagram
• the ‘ : |||| H — X ’ must be in a straight line [X being N, O or F]
• all partial charges and lone pairs shown
e.g. two molecules of water interacting:
how do you determine the shape of a molecule?
(1) group number of central atom + charge of molecule + number of bonds in the molecule
(2) ANS to (1) divided by 2 ( = number of electron pairs - this tells you the structure)
(3) ANS to (2) - number of bonds = number of lone pairs
(3) work out name of shape and bond angles (for every lone pair, -2.5°)
e.g. METHANE, tetrahedral, bond angles = 109.5° :
• central atom = C = group 4
• charge = 0
• 4 bonds in the molecule
• SO, tetrahedral, bond angles = 109.5°
name and bond angles of a molecule with a total of 2 electron pairs - 2 being bond pairs and 0 being lone pairs?
• linear
• 180°
e.g.
name and bond angles of a molecule with a total of 4 electron pairs - 2 being bond pairs and 2 being lone pairs?
• bent
• arrangement = tetrahedral as there’s a total of 4 pairs -> 109.5° - 2.5° - 2.5° [2 LP] = 104.5°
e.g.
name and bond angles of a molecule with a total of 3 electron pairs - 3 being bond pairs of electrons and 0 being lone pairs?
• trigonal planar
• 120°
e.g.
name and bond angles of a molecule with a total of 4 electron pairs - 3 being bond pairs of electrons and 1 being a lone pair?
• trigonal pyramidal
• (arrangement = tetrahedral as there’s a total of 4 pairs -> 109.5° - 2.5° [1 LP] =) 107°
name and bond angles of a molecule with a total of 4 electron pairs - 4 being bond pairs of electrons and 0 being lone pairs?
• tetrahedral
• 109.5°
e.g.
name and bond angles of a molecule with a total of 5 electron pairs - 5 being bond pairs of electrons and 0 being lone pairs?
• trigonal bipyramidal
• 90°
• 120°
e.g.
name of a molecule with a total of 5 electron pairs - 4 being bond pairs of electrons and 1 being a lone pair?
• seesaw
e.g.
name and bond angles of a molecule with a total of 5 electron pairs - 3 being bond pairs and 2 being lone pairs?
• T shaped
• 90°, 180°
e.g.
name and bond angles of a molecule with a total of 5 electron pairs - 2 being bond pairs of electrons and 3 being a lone pair?
• linear
• 180°
e.g.
name and bond angles of a molecule with a total of 6 electron pairs - 6 being bond pairs of electrons and 0 being lone pairs?
• octahedral
• 90°
e.g.
name and bond angle of a molecule with a total of 6 electron pairs - 5 being bond pairs of electrons and 1 being a lone pair?
• square pyramidal
• <90°
e.g.
name and bond angles of a molecule with a total of 6 electron pairs - 4 being bond pairs of electrons and 2 being lone pairs?
• square planar
• 90°
e.g.
factors that affect electronegativity?
• nuclear charge
• atomic radius
• shielding
trend in electronegativity down a group and why?
• decreases
• atomic radius and shielding increase (outweighs the increase in nuclear charge)
• weaker attraction between the nucleus and pair of electrons
trend in electronegativity across a period and why?
• increases
• atomic radius decreases (so atom gets smaller)
• nuclear charge increases
• same shielding
• stronger attraction between the nucleus and pair of electrons
when is a molecule polar?
• polar bonds present
• (the molecule isn’t symmetrical so) dipoles do not cancel each other out
why is ice less dense than water?
• water has hydrogen bonding
• there are spaces in the structure OR water molecules are held further apart OR water molecules are more spread out
why do macromolecules (such as diamond or silicone dioxide) have high melting points?
• giant covalent lattice held together by many strong covalent bonds
• lots of energy needed to break these bonds
how may a polar bond be drawn?
• partial charges
• arrow thing
when do polar bonds arise?
when the atoms have a significant difference in electronegativity
why are macromolecules (such as diamond and silicon dioxide) rigid?
many strong covalent bonds present which form a strong giant covalent lattice
why can graphite conduct electricity?
free/mobile delocalised electrons which are able to move through the structure
why do simple molecular compounds have low melting and boiling points?
• covalently bonded molecules held together by weak van der Waals forces
• not much energy required to overcome the vdW forces
why are simple covalent molecules poor conductors of electricity?
no charged particles present that are mobile and able to carry charge
even though it is a simple covalent molecule, why does water have an unusually higher boiling point?
• hydrogen bonding is present
• strongest intermolecular force
• more energy required to overcome these forces
why do differently shaped objects of titanium have similar strengths?
same metallic bonding OR same bond strength
why does a larger hydrocarbon have a higher boiling point than a smaller hydrocarbon?
larger one has stronger van der Waals’ forces between molecules as the molecules have a larger surface area and are packed closer together
Explain the polarity of a polar bond.
• ___ is more electronegative than ___
• So [more electronegative element] is δ- and [other one] is δ+
Explain the bond angles in a molecule where bond angles are equal.
(e.g. tetrahedral molecule - ALL bond angles are 109.5° - why?)
electron pairs repel equally
• B
• Na+ and F- have the same electron configuration
• Sodium ion has more protons than a fluoride ion, so has a greater attraction to the other electrons (so is the smaller ion)
How do lone pairs of a specific molecule influence the bond angle in it?
• Lone pairs repel more than bond pairs
• So, the bond angle will be lower
Why may molecules of the same substance NOT have dipole-dipole forces between them, despite each individual molecule having polar bonds?
e.g. CCl4
Symmetrical molecule, so dipoles cancel out
Types of bonding in ammonium chloride?
Covalent, dative covalent, and ionic
Suggest why magnesium is a liquid over a much greater temperature range compared to bromine.
Forces of attraction between the positive magnesium ions and delocalised electrons are stronger in liquid state
Explain how the value of the Cl-Al-Cl bond angle in AlCl3 changes, if at all, on formation of the compound H3NAlCl3.
• NH3 and AlCl3 form a dative covalent bond (electron pair donated from N of NH3 to Al) to produce H3NAlCl3
• Shapes of NH3 and AlCl3:
• Aluminium is now surrounded by 4 bonds, so H3NAlCl3 is tetrahedral
• So, Cl–Al–Cl bond angle decreases to 109.5° in H3NAlCl3
Use your knowledge of structure and bonding to explain why sodium bromide has a melting point that is higher than that of sodium, and higher than that of sodium iodide.
• Sodium has a giant metallic lattice structure and has metallic bonding – there is attraction between the positive ion and delocalised electrons
• NaBr and NaI have giant ionic lattice structures, and have ionic bonding – there is attraction between oppositely charged ions
• The ionic bonds are stronger than the metallic bonds
• Since a Br- ion is smaller than a I- ion, there is stronger attraction between the oppositely charged ions in NaBr than in NaI
By reference to the forces between molecules, explain why ammonia is very soluble in water.
hydrogen bonding between water and ammonia
