Topic 2a: bonding

ionic bond definition

electrostatic attraction between two oppositely charged ions which form when electron transfer from the outershell of one atom to another

cations and anions

cations: positive ions

anions: negative ions

ionic bond explaination (charge)

the larger the charge → stronger attraction

ionic bond explaination (ionic radius)

ionic radius depends on the number of shells and the number of protons in the nucleus

more shells → bigger ions

more protons → force of attraction on the electrons is stronger → smaller ions

isoelectronic ions

same electron configuration but different number of protons

properties of ionic compounds

because they contain giant lattices f ions strongly bonded together by electrostatic attraction

• hard, brittle

• high melting and boiling point

• soluble in water

• do not conduct electricity in solid but conduct in liquid (ions are free to move)

migration of ions experiment

using electrolysis and coloured ions (CuCr2O7) → orange negative ions move towards the positive electrode (Cr2O7 2-) and blue positive ions (Cu 2+) move towards the negative electrode

covalent bonding

strong electrostatic atraction between two nucleui and the shared pair of electrons

dative covalent bond

sharing of electrons is uneven → both shared electrons come from the same atom

example of dative covalent bond (Al2Cl6)

bond lengths and bond size of covalent bonds

bond length: shorter bonds are stronger → shared pair of electrons are closer to the oppositely charged nucleui → force of attraction is stronger

bond size: the smaller the atoms → stronger the bond → less shileding electrons between the nucleus and the outershell

metallic bonding

the strong electrostatic attraction between a lattice of positive metal ions and the sea of delocalised electrons

malleability and ductility of metals

malleable → can be hammered into shape

ductile → can be drawn out into a wire

conductivity of metals

can conduct heat in solid and liquid → because they have delocalised electrons that are free to move throughout the lattice

predicting the shapes of molecules

• counting the pairs of electrons around the central atom

(pairs of electron repel each other so they stay as far as way as possible to minimise electron repulsion)

2 electron pairs

linear, 180° bond angle,

3 electron pairs

• trigonal planar

• bent/v shaped

trigonal planar

120° bond angle

bent/ v shaped

2 bonding pair, 2 lone pair, 104.5° bond angle

4 electron pairs

• tetrahedral

• trigonal pyramidal

• bent

tetrahedral

4 bonding pair, 0 lone pair, 109.5° bond angle

trigonal pyramidal

3 bonding pair, 1 lone pair, 107° bond angle

5 electron pair

• trigonal bipyramidal

• 5 bonding pair, 0 lone pair, 120° and 90° bond angle

6 electron pairs

• octahedral

• 6 bonding pair, 0 lone pair, 90° bond angle

non bonding pair space taken up…

space taken up by non bonding pair (lone pair) is greater than space taken up by bonding pair, this is because lone pairs repel lone pairs more than lone pairs repel bonding pairs

(general rule: subtract 2.5 in bond angle for each lone pair present)

answer strucutre for if no lone pair or have lone pair present

no lone pair: the shape is ___ because there are __ bonding pairs and no lone pairs around the central atom which arrange themselves to a position of minimul repulsion

electronegativity

ability of an atom to attract the bonding electrons in a covalent bond

highest electronegativity

F has the highest electronegativity in the periodic table so it has the strongest attraction for those bonding electrons

electronegativity to tell types of bonding

ionic: large difference in electronegativity (>1.7)

ionic with covalent character: large electronegativity difference (>1)

polar covalent: different electronegativities

covalent: identical or very similar electronegatrivities

polar covalent bonds

usually bonded with highly electroegative atoms like F,O,N,Cl,S,P

polar molecules

uneven distribution of electrons caused by the difference in electronegativity within the bonded atoms

• eg. if the molecule is symmetrical and polar → the bond shape and the charge usually cancels out which makes it a non polar molecule

importance of intermolecular forces

• melting temperature

• boiling temperature

• density

• solubility

intermolecular forces explaination (strength)

the stronger the intermolecular forces → the more energy it requries to seperate molecules

(intermolecular forces are never strong tho)

3 types of intermolecular forces and strong or weak

1. london forces (weakest) → between all molecules

2. permanent dipole forces → only between polar molecules

3. hydrogen bonds (strongest) → only between molecules containing hydrogen bonded to NOF

permanant dipole - permanant dipole attraction

the slightly positive end on one molecule attracts the slightly negative end of another

instataneous dipole - instantaneous dipole attraction (london forces)

(neither contains a permanent dipole)

• at any moment, the distribution of electrons in a molecule could be momentarily uneven with greater electron density in one area → this produces instantaneous dipole

• the instantaneous dipole has the effect of inducing a second dipole on a neighboring molecule resulting in attraction between the molecules

(the more electrons in a molecule → more likely to have momentarily distributed unevenly electrons → stronger the london forces)

strength of london force…

depends on the number of electrons in a molecule

hydrogen bonding explaination

because when hydrogen is bonded, it only has one shell so it results in a poorly shielded proton → this means when its attracted to the lone pair of electrons on a electronegative atom of another molecule → produces strong [permanent dipole permanent dipole]

hydrogen bonding definition

electrostatic force of attraction between the poorly shielded proton of a hydrogen atom bonded to a small highly electronegative atoms such as NOF and a lone pair of electrons on a NOF atom of a neighbouring molecule

hydrogen bond bond angle

ALWAYS 180!!!!

water’s density is not 0…?

water is actually less dense below 4°C because of the open latticestructure of the ice

water’s boiling point is higher despite having less electrons??

this is because water forms hydrogen bonds and this makes its BP higher than others as it takes more energy to pull the water molecules away from their neighbours