Sem título

Página 1

Bonding & Structure

intermolecular forces

london (dispersion) forces (instantaneous dipole - induced dipole) · weakest type of imf · all atoms and molecules form these as they cause all atoms and molecules to be attracted to each other - because the electrons are constantly moving, the dipole are being created and destroyed all the time - even though dipole keep changing, the overall effect is that the atoms are attracted to each other · instantaneous (temporary) dipole = when electrons are likely to be more to one side of the atom · this dipole can induce another temporary dipole in the opposite direction of another atom - these two dipole are attracted to each other - the second dipole can induce another dipole in a third atom ect... · holds molecules in a molecular lattice arrangement by weak london forces · a.k.a = simple molecular structure properties · high melting points · if a solid has a stronger LF then more energy is required to overcome the imf · high boiling points - need to overcome the imf so that particles can escape - you need more energy to overcome stronger imf so if a liquid has a stronger london force, it has a higher boiling pts

· larger molecules have larger electron clouds, meaning stronger london forces - molecules with greater surface area have stronger london forces as they have a bigger exposed electron cloud alkanes · the longer the carbon chain, the stronger the LF as there's more molecular surface contact and more electrons to interact - as molecules get longer, it gets harder to separate them because it takes more energy to overcome LF

permanent dipole-permanent dipole · stronger than LF but weaker than H-bonding · polar molecules have these - the slightly positive and negative charges on polar molecules cause weak electrostatic forces of attraction between molecules · occur AS WELL AS london forces - molecules that form PD-PD as well as LF have high melting + boiling points

hydrogen bonding · strongest type of imf · only occurs when H bonds to F/O/N · F/O/N are highly electronegative so they draw in the bonding electrons away from the hydrogen atom - the bond is so polarised, and H has a high charge density as its so small, that H atoms form weak bonds with lone pairs of electrons on the F/O/N atoms - organic molecules that from H-bonds often contain -OH/-NH group high melting and boiling points - a lot of energy is required to overcome the imf ice floats on water - water molecules are arranged so that there is maximum number of hydrogen bonds - as ice melts, some of the hydrogen bonds are broken and the lattice breaks down - ice is much less dense than water so ice floats alcohols are less volatile (high boiling pts) than similar alkanes - all alcohols contain a polar hydroxyl group (-OH) that has a slightly negative charge on the O atom and a slightly positive charge on the H atom - this polar group helps to form h-bonds

solubility

in order to dissolve: · bonds in substance have to break · bonds in solvent have to break · new bonds have to form between substance and solvent polar solvents · made of polar molecules · but not all polar molecules can form hydrogen bonds non-polar solvents · such as hexane - bond to each other by london forces ionic substance · when ionic substances mix with water, the ions in ionic substance are attracted to the oppositely charged ends of the water molecules · hydration = ions are pulled away from the ionic lattice by water molecules which surround the ions alcohols · covalent but dissolve in water - because the polar O-H bond in alcohol is attracted to the polar O-H bond in water - hydrogen bonds from between the lone pairs on the slightly negative O and the slightly positive H · the more carbons there are, the less soluble the alcohol will be

halogenalkanes · contain polar bonds but their dipoles aren't strong enough to form hydrogen bonds with water · hydrogen bonds between water molecules is stronger than the bonds that would be formed with halogenalkanes so they don't dissolve · can form PD-PD bonds so they happily dissolve in polar solvents that also form PD-PD bonds non-polar substances · these have LF between their molecules and form similar bonds with non-polar solvents

intramolecular forces

ionic bonding · ionic bond = strong electrostatic attraction between two oppositely charged ions ionic charges - the greater the charge, the stronger the ionic bond and the higher the melting and boiling pts ionic radii - smaller ions can pack closer together than larger ions - electrostatic attraction gets weaker with distance so small, closely packed ions have stronger ionic bonding than larger ions - ionic compounds with small closely packed ions have higher melting and boiling points than ionic compounds made of large ions - ionic radius increases down a group as the atomic number (no. of protons) increases as extra electron shells are added - radii of isoelectronic ions decreases as the atomic number increases ions are positively or negatively charged atoms - ions are formed when electrons are transferred from one atom to another - cations = positive ions anions = negative ions - isoelectronic ions = ions of different atoms with the same number of electros ionic compounds form giant ionic lattice structures - each ion is electrosatically attracted in all directions to ions of the opposite charge properties · high melting point - ions held together by strong attraction · often soluble in water but not in polar solvents - particles are charged - ions are pulled apart by polar molecules but not by non-polar molecules · don't conduct electricity when solid - ions are in a fixed position by strong ionic bonds in a solid · conduct electricity when molten (liquid) or dissolved - ions are free to move and carry a charge in a liquid or solution · can't be shaped / brittle - repulsion between ions would be very strong so ionic compounds are brittle (they break when stretched)

covalent bonding giant covalent · diamond - each carbon atom is bonded to four neighbours in a tetrahedral arrangement · graphite - carbon atoms forms sheets, with each atom sharing three of its outer shell electrons with three other carbon atoms, leaving the fourth outer electron in each atom to move freely between the sheets · silicon dioxide - oxygen atoms between each silicon atom in a tetrahedral arrangement · properties · high melting points - a lot of energy is required to break bonds

· hard - very strong bonds through lattice arrangement

· good thermal conductors - vibrations travel easily through the stiff lattice

· insoluble - covalent bonds mean that the atoms are more attracted to their neighbours in the lattice than to solvent molecules

- don't contain ions

· conduct electricity? - NO = diamond + silicon dioxide

- no charged or free electrons as all the bonding electrons are held in localised covalent bonds

- YES = graphite

- the delocalised electrons are free to move along the sheet

· covalent bond = the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond · molecules are formed when two or more atoms bond together by covalent bonds dative covalent - where one atom donates both electrons to a bond / where both electrons come from one atom bond enthalpy - the positive nuclei are attracted to the area of electron density between the two nuclei (where the shared electrons are) · the two positively charged nuclei repel each other, as do the electrons · to maintain a covalent bond there has to be a balance between these forces · bond length = distance between the two nuclei is the distance where the attractive and repulsion forces balance each other · the higher the electron density between the nuclei, the stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length · the more electrons in the bond, the higher the electron density · more attraction = nuclei pulled closer together

metallic bonding form giant lattice structures · the electrons in the outermost shell of the metal atoms are delocalised (free to move) · this leaves a positive metal ion · overall lattice structure is made up of layers of positive metal ions separated by layers of electrons properties · high melting points - because of the strong metallic bonding with the number of delocalised electrons per atom - the more electrons, the stronger the bonding so the higher the melting point · malleable (can be shaped) - no bonds holding specific ions together and the layers of positive metal ions are separated by layers of electrons - ductile (can be sreched into a wire) = layers of metal ions can slide over each other without disrupting the attraction between the positive ions and the electrons · good thermal conductors - delocalised electrons can pass kinetic energy to each other · good electrical conductors - delocalised electrons are free to move and can carry a current · insoluble - because of the strength of the metallic bonds · metallic bonding = positive metal ions are electrostatically attracted to the delocalised negative electrons

electronegativity

· the ability of an atom to attract the bonding electrons in a covalent bond · the higher the electronegativity value, the more electronegative the element · fluorine is most electronegative (4.0) on the Pauling scale · most electronegative elements have higher nuclear charges (they have more protons) and smaller atomic radii · electronegativity increases across a period and decreases down a group (ignore the noble gasses) Pauling scale · if electronegativity value difference = 0 = purely covalent bond · if electronegativity value is more than 0.4 = polar

polarisation · in covalent bond, the bonding electrons sit in orbitals between two nuclei · if both atoms have similar or same electronegativities, the electrons will sit roughly midway between the two nuclei = non-polar · diatomic gases = non-polar as the atoms have equal electronegativities so the electrons are equally attracted to both nuclei · if bond is between two different electronegativitites , the bonding electrons will be pulled towards the more electronegative atom - this causes the electrons to be spread out so each atom has a partial charge - one is slightly positive whilst the other is slightly negative · in a polar bond, the difference in electronegativity between the two atom causes a dipole · dipole = a difference in charge between the two atoms caused by a shift in electron density in the bond · the greater the difference in electronegativity, the greater the shift in electron density, the more polar the bond · polar bonds don't always make polar molecules - a polar molecules has an overall dipole caused by the presence of a permanent change across the molecule - in a simple molecule - polar bond gives the whole molecule a permanent dipole - complex molecule (i.e: has double bonds) - if polar bonds are arranged so they point out in the opposite direction, they'll cancel out so molecule = NON-POLAR - if polar bonds point in the same direction - molecule = POLAR

shapes of molecules · shape depends on the number of pairs of electrons in the outer shell of the central atom · electron pairs repel each other as electrons are all negatively charged - the type of electron pair affects how much it repels the outer electrons o lone pairs repel more than bonding pairs - lone pairs repel more than bonding pairs - the greatest angles are between lone pairs of electrons o bond angles between bonding pairs are often reduced o bonding pairs are often pushed together by lone pair repulsion · lone pair - lone pair angles = biggest · lone pair - bonding pair angles = smaller than LP-LP but bigger than BP-BP · bonding pair - bonding pair angles = smallest