bonding

what is an ionic bond

electrostatic attraction between oppositely charged ions in a giant ionic lattice.

what is ionic bonding

lonic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

How are ions formed?

Atoms gain or lose electrons to achieve a stable noble gas electron configuration.

Cation = positive ion

Anion = negative ion

predicting charges on ions

Group 1 = 1+

Group 2 = 2+

Aluminium = 3+

D block

• Do not always achieve noble gas electron configuration.

• Can have various charges e.g iron

• Group 5 = 3-

• Group 6 = 2-

• Group 7 = 1-

what are the common molecular ions to know

Ammonium: NH⁺₄

Hydroxide: OH^-

Nitrate: NO^-₃

Hydrogen carbonate: HCO^-₃

Carbonate: CO^2-₃

Sulfate: SO^2-₄

Phosphate: PO^3-₄

rule for naming compounds

• Positive ion comes first and retains it's name

• Negative ions end in "ide"

• A negative ion with oxygen ends in "ate".

describe a giant ionic lattice

Oppositely charged ions are held in a regular 3D lattice by electrostatic attraction.

describe/explain the properties of ionic compounds

1. High melting and boiling points - lots of energy is needed to overcome the many strong electrostatic forces of attraction

between oppositely charged ions (can only be achieved at high temp)

2. Conduct electricity when molten or dissolved in solution. The ions are free to move around and carry a charge

3. Soluble in water - Water is polar. lon-dipole attractions form between water molecules and the ions, which can overcome the electrostatic attractions between the ions in the ionic lattice, causing the ions to separate and become Hydrated.

how to determine the strength of an ionic bond

• The strength of ionic bonds depends on charge densities of ions

• The smaller the ions + higher charge (greater charge density) = stronger ionic bond.

• E.g MgO has a higher melting point than NaCI because it has higher charges i.e Mg2+ and O2-.

• The magnesium ion is smaller than the sodium ion because magnesium it has more protons (12) that pulls it's electrons

more strongly

what is covalent bonding

a pair of two non metal atoms share some of their outer electrons to obtain a stable noble gas

electron configuration

how does sharing electrons hold atoms together?

atoms are held together by electrostatic attraction between the positive nuclei and the shared

pair of electrons.

There are forces of attraction and repulsion.

These balance when the nuclei are a particular distance apart (bond length).

how do atoms form covalent bonds?

• atoms use unpaired electrons in orbitals to form covalent bonds.

• The unpaired electron in the orbital of one atom is shared with an unpaired electron in an

orbital of another atom.

what is hybridisation

atoms can promote electrons into unoccupied orbitals on the same shell to form covalent bonds

what is a dative (coordinate bond)

one of the atoms provides both of the electrons which form the bond.

the atom that accepts the electron pair has an unoccupied orbital in its outer shell

the donating atom has a lone pair of electrons

describe/explain the physical properties of simple molecules

1. mp & bp

Low - intermolecular forces are weak, so little

energy is required to separate molecules

from each other

2. Electrical

Do not conduct electricity as they have no

ions or mobile electrons

3. Solubility

Tend to be more soluble in organic solvents

than in water; some are hydrolysed

depends on polarity and functional groups

describe essential features of Metallic bonding

• Metal atoms lose their outer shell electrons.

• These electrons form a "sea of delocalised electrons" which hold the positive ions in a lattice.

• The positive metal ions are arranged in a regular lattice surrounded by a sea of delocalised electrons.

• Metallic bonding involves electrostatic force of attraction between the positive metal ions and delocalised electrons.

• Atoms arranged in a regular close packed 3-dimensional crystal lattice.

The "sea of delocalised electrons" permeates the metal. The electron cloud binds the positive ions together.

describe what the strength of a metallic bond depend on

• charge on the ion = the greater the charge, and consequently the greater the number of delocalised electrons per ion, the stronger the electrostatic forces of attraction between the nuclei and the delocalised electrons.

• size of the ion = the smaller the ionic radius, the closer the delocalised electrons are to the positive nuclei, and the stronger electrostatic forces of attraction

• shielding = the more occupied electron shells, the more shielding, and the weaker the electrostatic forces of attraction between the positive nuclei and delocalised electrons.

describe/explain the physical properties of metals

1. High density = cations are tightly packed together.

2. High melting and boiling points = large amount of energy is needed to overcome the strong electrostatic forces of attraction between the cations and delocalised electrons.

3. Insoluble = electrostatic forces of attraction between metal atoms stronger than those between water molecules and metal atoms.

4. Excellent conductors of electricity and heat = Delocalised electrons are able to carry a charge and move through the lattice. They can also transfer thermal energy.

   • The more outer electrons a metal has the better it conducts.

5. Ductile and malleable = The layers of cations can slide over each other without disrupting the forces of attraction between the positive ions and delocalised electrons.

describe the trend of melting point across a period and down a group (for metals)

• Melting point for metals increases across the period. There are more delocalised electrons per atom and the ionic radius decreases.

• Melting point decreases down a group. lonic radius and shielding increase decreasing the attraction to the delocalised electrons.

what is melting point

The melting point is a measure of how strongly the delocalised electrons attract the cations.

what is electronegativity

the power of an atom to attract the pair of electrons in a covalent bond

describe and explain the Factors affecting electronegativity

1) Nuclear charge - more protons, stronger attraction

between nucleus and bonding pair of electrons.

2) Atomic radius - closer to the nucleus, stronger attraction between nucleus and bonding pair of electrons.

3) Shielding - less shells of electrons between the nucleus and the electrons, less shielding (less repulsion), stronger attraction between nucleus and bonding pair of electrons.

describe the trend of electronegativity down a group

Electronegativity decreases

• Atomic radius increases

• More shielding

• Less electrostatic force of attraction between nucleus and bonding pair of electrons

describe the trend of electronegativity across a period

Electronegativity increases

• Atomic radius decreases

• More nuclear charge

• Stronger attraction between nucleus and bonding pair of electrons

what is the condition to have a polar bond

uneven distrubition of electrons

Summarise the spectrum of chemical bonding from pure covalent to pure ionic in

terms of electron location.

pure covalent = electrons shared equally

polar covalent = electrons not shared equally

polar ionic = distorted ions i.e a high charged cation attracts the outer electrons of an anion causing the electron cloud to stretch to the point where it is in between the 2 nuclei and it looks like sharing of electrons so it has covalent character.

pure ionic = electrons all on one atom

polarised ionic bonds are favoured by?

positive ions with high charge density i.e small ion and high charge.

how is a dipole formed?

• The Cause: Electrons in a molecule or atom can move from one end to another.

• The Effect: This creates an uneven distribution of charge, causing one side to be slightly positive and another side to be slightly negative.

• The Result: This atom or molecule becomes a temporary dipole.

Explain how an instantaneous dipole in one non-polar atom or molecule induces a dipole in a neighbouring atom or molecule. (4 marks)

• The slightly negative (8-) end of the first particle (Atom A) interacts with a neighbouring particle (Atom B).

• The excess negative charge on Atom A repels the electron cloud of Atom B, pushing its electrons towards the far side.

• This creates an uneven electron distribution in Atom B, inducing a temporary dipole in it.

• This induced dipole is temporary and only exists while the particles remain in close proximity.

Describe how London dispersion forces (instantaneous dipole-induced dipole interactions) are formed and broken between two non-polar atoms or molecules. (3 marks)

• The slightly negative charge on one atom attracts the slightly positive charge on an adjacent atom.

• A temporary force of attraction/temporary dipole-dipole interaction is formed between them

• As the atoms/molecules move apart, the electron distribution changes and the dipole interaction is destroyed.

what are the types of intermolecular forces and rank their weakness

Hydrogen Bonding = weak

Permanent Dipole-Dipole = weaker

Van der Waals (induced dipole-dipole) = weakest

what is a permanent dipole

A permanent dipole is a charge imbalance across a bond or molecule that is always present.

this uneven distrubition of charge is due to difference in electronegativity

Explain how permanent dipole-dipole forces arise between two polar molecules.

• The molecule has a permanent uneven distribution of charge because a more electronegative atom pulls shared electrons toward itself.

• This creates permanent slightly positive and slightly negative charges on opposite sides of the molecule.

• The adjacent molecule (Molecule B) also possesses its own permanent dipole.

• The slightly negative end of Molecule A aligns with and attracts the slightly positive end of Molecule B.

what does a higher difference in electronegativity mean?

A higher difference in electronegativity directly makes a bond more polar

when does a molecule have polar bonds

when there is a difference in electronegativity between the atoms in the molecule.

when is a molecule polar

A molecule is polar if it has polar bonds AND those bonds are arranged non-symmetrically i.e in the molecule there is clear side that is mostly positive and another clear side that is mostly negative so their dipoles don't cancel out.

State the rule regarding which types of intermolecular forces are present in non-polar

molecules, polar molecules without hydrogen bonding, and molecules that exhibit

hydrogen bonding. (3 marks)

• Any molecule with van der Waals forces only has van der Waals forces.

• Any molecule with permanent dipole- dipole forces will have both permanent dipole-dipole forces and van der Waals forces.

• Any molecule with hydrogen bonding will have all three: hydrogen bonding, permanent dipole-dipole forces, and van der Waals forces.

what real life way is there to see if a molecule is polar

test to see if a charged rod would deflect the molecule as it is rushing from a burette

describe hydrogen bonding

• Hydrogen bonding is a particularly strong dipole-dipole interaction.

• It occurs in molecules with F-H, O-H and N-H bonds, because F, O and N are the three most

electronegative elements.

• The polar bond leaves the H nucleus exposed as H only has one electron.

• Hydrogen bonds occur between a 8+ H atom of one molecule and the lone pair of electrons an O, N or F atom of another molecule.

• This is simply a strong intermolecular force - it is NOT a bond!

what are the rules for drawing hydrogen bonds

draw hydrogen bonds between water molecules

if we have 2 non polar molecules so they just have van der waal forces how would you able to tell which molecule has the higher boiling point

• the bigger molecule (i.e that has more electrons) will have a higher boiling point

Term (Question): Describe and explain the general trend in boiling points of hydrides

down a group.

Trend: Boiling point increases down a group.

Explanation: Larger molecules have more electrons → stronger London dispersion

forces (Van der Waals forces) → more thermal energy required to overcome them.

Term (Question): Describe and explain the general trend in boiling points of covalent hydrides across a period (from left to right, excluding hydrogen-bonding hydrides).

Trend: Boiling point increases as the number of hydrogen atoms decreases (moving

left to right).

Explanation: Central atom electronegativity increases → molecular polarity

increases → stronger dipole-dipole attractions between molecules.

Explain why the boiling points of HF, H₂O, and NH3 do not follow the general trend of increasing down their respective groups.

Answer: These specific hydrides have anomalously high boiling points because their molecules can form hydrogen bonds (due to the high electronegativity of F, O, and N bonded to H), which require significantly more thermal energy to break than simple London dispersion forces.

describe the structure of ice and why it is less dense than water?

• Water molecules possess induced dipole-dipole, permanent dipole- dipole and hydrogen bonds.

• In the liquid state, the hydrogen bonds break and reform relatively easily and the molecules move about.

• The hydrogen bonding gives liquid water a surface tension, used by pond skaters

• when water freezes the molecules stop moving about and optimise their hydrogen bonds, leading to a crystaline solid.

• In this state the water molecules are slightly less close packed than in liquid so ice is less dense than water