Module 2 - Atoms, Ions and Compounds

The nuclear atom

• consists of a nucleus made up of two types of subatomic particles - protons and neutrons

• a 3rd type of subatomic particle, called an electron, occupies a region outside the nucleus

• electrons are arranged around the nucleus in shells

Mass of the subatomic particles

• proton - 1

• neutron - 1

• electron - 1/1836

Charges of the subatomic particles

• proton - +1

• neutron - 0

• electron - -1

• the total positive charge from protons is cancelled by the total negative charge from electrons

• the overall charge of an atom is 0 - an atom is neutral

Structure of the atom

• nearly all of an atom’s mass is in the nucleus

• atoms contain the same number of protons as electrons

• most atoms contain the same number of, or slightly more, neutrons than protons

• as the nucleus gets larger, more and more neutrons are needed

Atomic number - the identity of an element

the number of protons in the nucleus of an atom

Isotope

• atoms of the same element with different numbers of neutrons and different masses

• most elements are made up of a mixture of isotopes

Isotopes and chemical reactions

• different isotopes of the same element have the same number of electrons

• the number of neutrons has no effect on reactions of an element

• different isotopes of an element therefore react in the same way

• there may be small differences in physical properties - with higher mass isotopes of an element having a higher melting or boiling point or density - but the chemical reactions are the same

Atomic structure of ions

• an ion is a charged atom - the number of electrons is different from the number of protons

• cations (+ve) are atoms with fewer electrons than protons

• anions (-ve) are atoms with more electrons than protons

Relative isotopic mass

the mass of an isotope relative to 1/12th of the mass of an atom of carbon-12

Relative atomic mass

• the weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12

• it takes account of the percentage abundance of each isotope and the relative isotopic mass

Determination of relative atomic mass (mass spec)

1. A sample is placed in the mass spectrometer

2. The sample is vaporised and then ionised to form positive ions

3. The ions are accelerated. Heavier ions move more slowly and are more difficult to deflect than lighter ions, so the ions of each isotope are separated

4. The ions are detected on a mass spectrum as a mass to charge ratio (m/z). Each ion reaching the detector adds to the signal, so the greater the abundance, the larger the signal

5. m/z = relative mass of ion / relative charge on ion

Simple ions

• atoms of metals on the left of the periodic table lose electrons to form cations

• atoms of non-metals on the right of the periodic table gain electrons to form anions

Polyatomic ions

Avogadro’s constant

6.02×10²3

Molar mass

mass per mole of a substance

Mole

one mole is the amount of a substance that contains 6.02×10²3 atoms

Molecular formulae

• the number of atoms of each element in a molecule

• e.g H2, O2, N2 etc

Empirical formulae

simplest whole number ratio of atoms of each element in a compound

Relative molecular mass

• compares the mass of a molecule with the mass of an atom of carbon-12

• can be calculated by adding together the relative atomic masses of the elements making up a molecule

Relative formula mass

• compares the mass of a formula unit with the mass of an atom of carbon-12

• it is calculated by adding together the relative atomic masses of the elements in the empirical formulas

Hydrated salts

• many coloured crystals are hydrated - water molecules are part of their crystalline structure

• this water is known as water of crystallisation

• when blue crystals of hydrated copper sulfate are heated, bonds holding the water within the crystal are broken and the water is driven off, leaving behind white anhydrous copper sulfate

• without water, the crystalline structure is lost and a white powder remains

Concentration of a solution

the amount of solute, in moles, dissolved in each 1dm³

Standard solution

a solution of known concentration

Molar gas volume

• the volume per mole of gas molecules at a stated temperature and pressure

• at RTP the molar gas volume = 24.0 dm³ mol-1

Assumptions for molecules making up an ideal gas

• random motion

• elastic collisions

• negligible size

• no intermolecular forces

Ideal gas equation

Conversions for ideal gas equation

Stoichiometry

• in a balanced equation, the balancing numbers give the ratio of the amount, in moles, of each substance

• this ratio is called the stoichiometry of the reaction

What do chemists use balanced equations for?

• the quantities of reactants required to prepare a required quantity of a product

• the quantities of products that should be formed from certain quantities of reactants

• these quantities can then be changed to adjust the scale of a preparation

Theoretical yield

the maximum possible amount of product

Why is theoretical yield difficult to achieve?

• the reaction may not have gone to completion

• side reactions may have taken place alongside the main reaction

• purification of the product may result in loss of some product

• therefore the actual yield obtained from a reaction is usually lower than the theoretical

Percentage yield

actual yield / theoretical X 100

Limiting reagent

• a reactant that is totally consumed when the chemical reaction is completed

• the amount of product formed is limited by this reagent, since the reaction cannot continue without it

Atom economy

• a measure of how well atoms have been utilised

• sum of molar masses of desired products / sum of molar masses of all products X 100

• atom economy is based solely on the balanced chemical equation for a reaction and assumes a 100% yield

Reactions with high atom economies

• produce a large proportion of desired products and few unwanted waste products

• are important for sustainability as they make the best use of natural resources

Acid

when dissolved in water, releases hydrogen ions as protons (H+), into the solution

Strong acid

completely dissociates in aqueous solution e.g hydrochloric acid

Weak acid

partially dissociates in aqueous solution e.g ethanoic acid

Base

• neutralised an acid to form a salt

• e.g metal oxides, metal hydroxides, metal carbonates and ammonia

• an alkali is a base that dissolves in water releasing hydroxide ions into the solution

Neutralisation

• H+(aq) ions react with a base to form a salt and neutral water

• the H+ ions from the acid are replaced by metal or ammonium ions from the base

• metal oxides and alkalis neutralise acids to form salt and water only

• metal carbonates form salt, water and carbon dioxide

Titration

• a technique used to accurately measure the volume of one solution that reacts exactly with another solution

• they can be used for - finding the concentration of a solution, identification of unknown chemicals and finding the purity of a substance

Oxidation number

• based on a set of rules that apply to atoms, and can be though of as the number of electrons involved in bonding to a different element

• use of oxidation numbers helps when writing formulae and balancing electrons as a check that all electrons have been accounted for

Oxidation rules for elements

• oxidation number for elements is always 0

• in a pure element, any bonding is to atoms of the same elements

Oxidation rules for compounds and ions

Working out oxidation numbers

sum of the oxidation numbers = total charge

Redox reactions

• involve reduction and oxidation

• if one process happens so must the other

Reduction and oxidation in terms of oxygen

• oxidation is addition of oxygen

• reduction is removal of oxygen

Redox in terms of electrons

• reduction is the gain of electrons

• oxidation is the loss of electrons

Redox in terms of oxidation number

• reduction is a decrease in oxidation number

• oxidation is an increase in oxidation number