Chemistry : Chemical equilibria AS level

reversible reactions

products can react to reform the original reactants

dynamic equilibrium

a state where the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant concentrations of products and reactants.

A closed system

in which no matter is exchanged with the surroundings, allowing equilibrium to be established.

Conditions for equlibrium of gases

equilibrium can onlybe reached in a closed system

Le Chatelier's Principle

states that if a dynamic equilibrium is disturbed, the system will adjust to counteract the change and restore equilibrium.

position of the equilibrium

refers to the relative amounts of reactants and products at equilibrium, which can be influenced by changes in concentration, temperature, and pressure.

Increase in concentration of a reactant

Equilibrium shifts to the right , (reduce the effect of an increase in the concentration of a reactant)

Decrease in concentration of a reactant

Equilibrium shifts to the left

Effects of pressure

Changes in pressure only affect reactions where the reactants or products are gases

Increase in pressure causes the equilibrium to shift towards the side with fewer moles of gas, reducing the overall pressure.

Decrease in pressure causes the equilibrium to shift towards the side with more moles of gas, increasing the overall pressure.

equilibrium expression

Expression that links equilibrium constant to the concentrations of reactants and products at equilibrium. Kc only changes if the temperature of reaction changes

Solids are ignored in equilibrium expressions

Mole fraction

Partial pressure

equals mole fraction times the total pressure

Comparing Kc and Kp

Kc is the eqilibrium constant based on concentrations, while Kp is based on the partial pressure of gases

Acids

substances that release hydrogen ions when theydissolve in water

Acid + Base is a neutralization reaction. It forms salt and water

Base

a substance that accepts hydrogen ions or a compound that contains oxide or hydroxide ions.

The Brønsted-Lowry Theory

defines acids and bases in terms of proton transfer between chemical compounds

The Brønsted-Lowry Theory Acid/Base

species that gives away a proton (for acids )

species that accepts a proton using its lone pair of electrons (for bases)

GK

Species that can act both as acids and bases are called amphoteric

Strong Acids

an acid that dissociates almost completely in aqueous solutions. F.E: HCl, nitric acid, sulfuric acid

Electronegativity

is the ability of an atom to attract a a pair of electrons towards itself in a covalent bond

Pauling scale

used to assign the value of e;lectronegativity

What happens when there is an increase in no.of protons

causes the nuclear attraction to increase as well as increasing electronegativity - therefore larger nuclear attraction which bonds the electrons more strongly

atomic radius

is the distance between the nucleus and the electrons in the outer most shell

Sheilding effect

grearter the no.of shells, less the outer electron is attracted to it s nucleus

trends in electronegativity - down a group

Going down the group there is a decrease on electronegativity - as sheilding and atomic radius increase down a group

trends in electronegativity - across a period

across a period electronegativity increases - nuclear charge across a period increases whereas the shielding remains the same. Atomic radius decreases

Covalent bonds

are formed by sharing a pair of electrons between two atoms

• occurs between 2 non-metals

• covalent compounds are crystalline lattice

Diatomic molecules

have equal distribution of bond pair leads to non polar molecule

laws of electronegativity

the least electronegative atom’s electron will transfer to the other atom - leading to an ionic bond

Ionic bond

involves the transfer of electrons froma metal to a non metal leaving it with a full outer shell

Electrostatic attractions

are formed between the oppositely charged ions to form ionic compounds - this form of attraction is very strong and requires a lot of energy to overcome

How are atoms affected by the delocalised electrons ?

Metal atoms become positively charged

double covalent bonding

triple covaent bonding

dative covalent bonding or coordinate bonding

iswhen the both electrons are from the same atom - for example

• Ammonium ion

• Aluminum chloride

Pressure

• volume is inversley proportional to the pressure

• temparature is directly proportional to pressure

Kinetic theory of gases

• gas molecules move very fast and randomly

• hardly have any volume

• gas molecules do not attract or repel each other

• have elastic collisions

Ideal gases

Ideal gases depend on its

• pressure

• temperature

Volume is directly proportional to the temperature

Limitations of ideal gas law

Ideal gases do not obey the kinetic theory of gases at very high or low temperatures, because

• molecules are close to each other

• volume of gases isn’t negligible

• pd-pd forces between the molecules

• attractive forces pull molecules away from container walls

ideal gas equation

pV = nRT

• p = pressure in pascals

• R = gas constant (8.31 j/k mol

• n = number of moles of gas

• T = temperature in kelvin (K)

Giant Ionic lattices

Type of lattice formed depends on the sizes of the anion and the cation. For example:

• MgO and NaCl are cubic

Covalent lattices

Covalent compunds can either be arranged in simple molecular or giant molecular

Example sof simple molecular:

• Iodine, Ice, ,buckminsterfullerene

Examples of giant molecular:

• graphite, diamond and silicon(IV) oxide

giant metalllic lattices

• in which metal ions are surrounded by a sea of delocalised electrons

• often packed in hexagonal layers or in a cubic arrangement

Physical Properties of Ionic bonding & giant ionic lattice structures

• Ionic compunds are strong

• They are brittle

• Have high melting and boiling points (depends on its charge density)

• soluble in water

• can form ion-dipole bonds

• conduct electricity when molten or in solution

Metallic bonding & giant metallic lattice structures

• Metallic compunds are malleable

• strong and hard

• high melting and boiling points

• Pure metals are insoluble in water

• conduct electricity while in solid or liquid state

Covalent bonding and simple covalent lattice

• have low melting and boiling points

• most compounds are insoluble in water

• do not conduct electricity

Covalent bonding & giant covalent lattice structures

• high melting and boiling points

• insoluble in water

• only graphite conducts electricity

• can be hard or soft (graphite soft, diamond hard)

Standard enthalpy change of reaction(endo and exo)

The enthalpy change when the reactants in the stoichiometric equation react to form the products under standard conditions

Standard enthalpy change of formation(endo and exo)

The enthalpy change when 1 mole of compound is formd from its elements USC

Standard enthalpy change of combustion(exo)

The enthaply change when one mole of substance is burnt in excess oxygen USC

Standard enethalpy change of neutralisation(exo)

The enthalpy change when one mole of water is formed by reacting and acid and an alkali USC

Exothermic reactions

• Heat is given off by the reaction to the surrounding

• products have less energy than the reactants

• exothermic reactions are thermodynamically possible

Endothermic reactions

• Heat is absorbed by the system from the surroundings

• Enthalpy change of endo is always positive

• products have more energy than the reactants

reaction pathway diagrams

activation energy (Ea)

The minimum energy vthat colliding particles must have for a collision to be successful and a reaction to take place

standard conditions

• pressure - 101kPa

• temperature - 298K

Bond breaking

• Energy need to overcome the attractive forces so its endothermic

Bond forming

• Energy is released from the reaction to the surroundings when new bonds are formed, therefore its exothermic

Calorimetry

a technique used to measure changes in enthalpy of chemical reactions

specic heat capacity

• The energy needed to increase the temperature of 1g of a substance by 1 degree celcius

SHC - equation

q = mcT

• q = heat transferred, J

• m = mass of the water, g

• c = specific heat capacity

• T = temperature change

Rate of reaction

• the speed at which a chemical reaction takes place

Collision theory

For a chemical reaction to occur

• the particles must collide with each other in the correct orientation

• must colide with enough energy

Collision frequeny

number of collisions per unit time