Chapter 6 (shapes of molecules + intermolecular forces)

• 0 lone pairs of electrons leads to a bond angle of 109.5 degrees (with exceptions) each lone pair of electrons decreases the bond angle by 2.5 degrees

• BF3, 3 atoms surrounding the central atom, with a bond angle of 120 degrees

• CO2 since there are no lone pairs of electrons, the bond angle is 180 degrees

• CH4, no lone pairs and 4 bonded atoms pairs, repel each other equally, resulting to a bond angle of 109.5 degrees

• NH3, 1 lone pair of electrons and 3 bonded atoms, therefore bond angle of 107 degrees

• H20, the central atom has 2 lone pairs of electrons, resulting in a bond angle of 104.5 degrees

• SF6, with 6 bonded atoms surrounding the central atom, with no lone pairs of electrons, resulting in a bond angle of 90 degrees

• PF5, 5 bonding pairs of electrons, no lone pairs, however this an exception where 3 atoms have a bond angle of 90 degrees and 2 atoms have a bond angle of 120 degrees

• The valence shell electron pair repulsion theory (VSEPR) predicts the shape and bond angles of molecules • In a molecule, the bonding pairs of electrons will repel other electrons around the central atom forcing the molecule to adopt a shape in which these repulsive forces are minimised

• Electron pairs repel each other as they have the same charge • Lone pair electrons repel each other more than bonded pairs • Repulsion between multiple and single bonds is treated the same as for repulsion between single bonds • Repulsion between pairs of double bonds are greater • The most stable shape is adopted to minimize the repulsion forces

• electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself • this means that the electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical • This is because of the ability of the positive nucleus to attract negativity charged electrons, towards it • the Pauling scale is used to assign a value of electronegativity

• Nuclear charge • Atomic radius • shielding

• Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom • An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells • Therefore, an increased nuclear charge results in an increased electronegativity

• The atomic radius is the distance between the nucleus and electrons in the outermost shell • Electrons closer to the nucleus are more strongly attracted towards its positive nucleus • Those electrons further away from the nucleus are less strongly attracted towards the nucleus • Therefore, an increased atomic radius results in a decreased electronegativity

• Filled energy levels can shield (mask) the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus • Sodium (period 3, group 1) has higher electronegativity than caesium (period 6, group 1) as it has fewer shells and therefore the outer electrons experience less shielding than in caesium • Thus, an increased number of inner shells and subshells will result in a decreased electronegativity

• there is a decrease in electronegativity going down the group • The nuclear charge increases as more protons are being added to the nucleus • However, each element has an extra filled electron shell, which increases shielding • The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii • Overall, there is decrease in attraction between the nucleus and outer bonding electrons

• Electronegativity increases across a period • The nuclear charge increases with the addition of protons to the nucleus • Shielding remains relatively constant across the period as no new shells are being added to the atoms • The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period of the periodic table • This results in smaller atomic radii

• When 2 atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar • When 2 atoms in a covalent bond have different electronegativity the covalent bond is polar and the electrons will be drawn towards the more electronegative atom

• the electron distribution is asymmetric • the less electronegative atom gets a partial positive charge (delta positive) • the more electronegative atom gets a partial negative charge (delta negative) • the greater the difference in electronegativity the more polar the bond becomes

• The dipole moment is a measure of how polar a bond is • The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole

• some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such a way that the individual dipole moments cancel eachother out e.g CCl4

• Induced dipole – dipole forces are also called London dispersion forces or van der Waals’ forces • Permanent dipole – dipole forces (also called van der Waals’ forces) are the attractive forces between two neighbouring molecules with a permanent dipole • Hydrogen Bonding are a special type of permanent dipole - permanent dipole forces • Intramolecular forces are stronger than intermolecular forces For example, a hydrogen bond is about one tenth the strength of a covalent bond

• The electron charge cloud in non-polar molecules or atoms are constantly moving • During this movement, the electron charge cloud can be more on one side of the atom or molecule than the other • Because the electron clouds are moving constantly, the dipoles are only temporary • Therefore the greater the number of electrons the molecule has, the stronger the induced dipole-dipole forces

• Polar molecules have permanent dipoles • The molecule will always have a negatively and positively charged end • permanent dipole forces are stronger than induced dipole forces