Chemistry - Chapter 5: Chemical Bonding

electronegativity

the relative power of attraction an atom of an element has for the shared pair of electrons in an atom in a single covalent bond

metals vs non-metals in electronegativity

metals generally have low electronegativity values as they are electropositive, while non-metals have relatively high electronegativity values

octet rule

when bonding occurs, atoms tend to gain, lose or share electrons when they combine with other atoms in order to gain a stable octet (8) of 8 electrons in their valence shell, which corresponds to the electron configuration of the nearest noble gas

exceptions to the octet rule

• transition metals - often have more than 8 electrons in their valence shell

• elements near helium, which is stable with only 2 outer shell electrons

cation

positively charged ion, formed when an atom loses electrons & no. of protons remain the same

anion

negatively charged electron, formed when an atom gains electrons

ionic bond

the force of attraction between oppositely charged ions - e.g. Na+ and Cl- = NaCl

it is formed due to the electrostatic attraction between oppositely charged ions caused by the complete transfer of electrons from one atom to another

remember, the particles in ionic bonding are IONS

so don’t talk abt atoms + use the dot and cross diagrams

properties of ionic compounds

• generally solid crystalline compounds at room temperature due to individual ions attracting other ions (extremely strong forces of attraction)

• high melting and boiling points due to a lot of energy being needed to break the extremely strong forces of attraction between the ions

• very soluble in polar solvents such as water/H2O, as the attraction between the ions in the compound and polar water are strong enough to pull the crystalline lattice apart

• conduct electricity when molten or dissolved state as ions are free to move

• undergo fast reactions

• hard and brittle

• occurs between a metal and non-metal

shape of ionic compound

• 3D crystal lattice shape which is surrounded by others of opposite charge

• this is because it does not result in the formation of individual molecules

why are ionic compounds crystalline?

because they are made up of an orderly arrangement of oppositely charged ions which give rise to a lattic structure

why are ionic compounds solid?

there are extremely strong attractions between the oppositely charged ions which hold the ions tightly together, giving a solid structure

why do ionic compounds have high boiling and melting points

the extremely strong attractions between the oppositely charged ions require large amounts of energy to break, hence ionic compounds have high b.p and m.p

valency

number of electrons an atom will lose, gain or share in order to have the same electronic configuration of the nearest noble gas

covalent bonding

sharing of electrons between atoms instead of losing or gaining electrons - occurs between non-metals

here, electrons are normally localised (not free to move from one atom to the next, but are held between the two atoms)

single covalent bond

one pair of electrons shared

double covalent bond

when atoms share 2 pairs of electrons + shorter but stronger than a single covalent bond

triple covalent bond

when atoms share 3 pairs of electrons + stronger than a single or double covalent bond and is shorter than a double covalent bond

sigma bonding

head-on / end-on overlap of orbitals

3 ways: head on, s-orbital overlapping with p-orbital, 2 p-orbitals overlapping head on

1 sigma bond is one single bond

occurs with s-s, s-p, p-p orbitals

stronger than a pi bond

happens before a pi bond as there is a greater/stronger overlap and will give stability

*all single bonds are sigma bonds

pi bonding

sideways overlap of p-orbitals

1 pi bond has 2 points of overlap - top and bottom of p-orbital

happens after a sigma bond

weaker than a sigma bond

double bond

sigma + pi bond

triple bond

sigma + 2 pi bonds

properties of covalent bonds

• low melting and boiling points

• generally gases, liquids of soft solids at room temperature

• don’t conduct electricity due to no ions

• undergo slower reactions

• occurs between 2 non-metals

electronegativity differences to predict bonding

• 0: pure covalent

• 0 > 0.5: slightly polar covalent

• <1.7: polar covalent

• >1.7: ionic

valence shell electron pair repulsion theory

• electron pairs repel each other so that they are as far apart as possible

• not all electron pairs have the same repulsion

linear

• 2 bond pairs

• 0 lone pairs

• 2 valence electrons

• 180 degree angle

• e.g. CO2, HCl, CN, BeH2

trigonal planar

• 3 bond pairs

• 0 lone pairs

• 3 valence electrons

• 120 degree angle

• e.g. BH3, AlCl3, BCl3

tetrahedral

• 4 bond pairs

• 0 lone pairs

• 4 valence electrons

• 109.5 degree angle

• e.g. CH4, SiCl4

pyramidal

• 3 bond pairs

• 1 lone pair

• 3 valence electrons

• 107 degree angle

• e.g. NH3, PCl3

planar v-shaped

• 2 bonding pairs

• 2 lone pairs

• 2 valence electrons

• 104.5 degree angle

• e.g. H2O, OF2, SBr2

intramolecular bonding

bonding within molecules

e.g. ionic and covalent bonding

intermolecular bonding

bonding between molecules

e.g. Van der Waals forces, dipole-dipole attractions, hydrogen bonds

hydrogen bonding

the force of attraction between molecules when a hydrogen, which is covalently bonded to a smaller, more electronegative element

only when hydrogen is bonded to F, O, N

• they are relatively weak but exert a large influence on the physical and chemical properties on the molecules that form them

• due to their presence, more energy is needed to break them, hence water has a higher b.p compared to h2s

• also explains that h2o is liquid at room temp. whereas h2s is a gas - due to the intermolecular forces of attraction between h2s molecules are considerably weaker than h2o, so they are freer to move

• h-bonding also accounts for the high solubilities of polar compounds in water

dipole-dipole attractions

a force of attraction between the opposite ends of permanent dipoles on neighbouring molecules - a pair of opposite charges

→ these are the weak intermolecular forces of attraction between polar molecules. they are described as being permanent dipoles as the molecules are constantly polar

→ the energy needed to break them varies (more than that needed to break VdW & less needed to break Hb)

dipole

forces of attraction between the negative pole of one polar molecule and the positive pole of another polar molecule

temporary dipoles

will only exist for an instant

Van der Waals forces

a force of attraction between temporary instantaneous dipoles induced on neighbouring non-polar molecules

they are weak attractive forces between molecules, caused by the formation of temporary dipoles in non-polar molecules

they are the only forces of attraction between non-polar molecules (pure covalent)

• these are the weak intermolecular forces of attraction that exist between all states of matter

• caused by internal shifts in the electrons in the molecule, hence a temporary dipole is set up in the molecule

• only intermolecular forces between gaseous molecules

-ate/-ite

-ide

-contains oxygen + elements in the name

-contains only elements in the name

explain why the energy to break a C≡C bond is not 3xx the energy to break a C-C

• pi bond is weaker than sigma bond, so we don’t need 3x as much energy

• sigma form first, but pi are broken first

types of covalent bonding

Pure covalent bonding

• defn: occurs when electrons are shared equally between atoms

• temporary dipoles in the molecule

• extremely weak dipole-dipole bonds, so Van der Waal’s forces between molecules

• remember that bigger molecules create greater surface area, so more surface area = more Van der Waal’s forces between molecules (more energy needed to break more Van der Waal’s forces

Polar covalent bonding

• defn: occurs when electrons are shared unequally between atoms, so one atom has a slightly more negative charge than the other

• permanent dipoles

• this leads to dipole-dipole attractions // hydrogen bonding (w/ h2o, nh3, hf)

Ionic bonding

• permanent dipoles

• ionic attractions

Dative covalent bonding (never asked)

polar molecule

different centres of positive and negative charge and have dipole moments - electrons in covalent bonds shared unequally

so

the slight positive and slight negative poles of a molecule do not coincide

• can be proved w/ charged comb vs stream of water

• also use the shapes of molecules for these

boiling points and states of matter

determined by strength of intermolecular forces of attraction & size of the molecule

the strong/heavier the molecule, the more energy required to break them

• ionic compounds are extremely polar and have the strongest forces of attraction, hence high bp and usually crystalline solids

• polar covalent compounds that have hydrogen bonding between the molecules have relatively lower b.p and may be liquids

• polar covalent compounds that have dipole-dipole interactions have bp that are slightly weaker and may be gases

• compounds with the very weak VdW’s forces have extremely low bp and generally tend to be gases, but exceptions occur