Chemistry F321 Atoms, Bonds and Groups

Proton - Mass = 1, Charge = +1, Location = Nucleus

Neutron- Mass = 1, Charhe = 0, Location = Nucleus

Electron - Mass = 1/2000, Charge = -1, Location = Orbitals

• The nucleus contains almost all of the mass of an atom because that is where the protons and neutrons are found
• The nucleus of an atom contains all of the positive charge
• The electrons are outside the nucleus and, therefore, so is the negative charge
• Atomic number              = number of protons in the nucleus
• Mass number                 = number of protons and neutrons in the nucleus
• Number of neutrons       = mass number – atomic number
• Number of electrons       = number of protons (in a neutral atom)

• Isotopes are atoms with the same number of protons but different numbers of neutrons (and different masses)
• For example, chlorine has two isotopes ³⁵Cl and ³⁷Cl. Both have 17 protons but they have 18 and 20 neutrons, respectively
• Isotopes of an element have the same chemical properties because they have the same electron arrangement
• Ions are formed when atoms gain or lose electrons.
• As an atom Cl has 17 electrons. A Cl^- ion has gained one electron so it now has 18.
• As an atom Na has 11 electrons. A Na⁺ ion has lost one electron so it now has 10.

Relative atomic mass is the average mass of an atom of an element taking the mixture of isotopes into account.

Example:

75% of Cl atoms have a mass number of 35

25% of Cl atoms have a mass number of 37

Average mass of a Cl atom         = (mass no x percent/100) + (mass no x percent/100)

                                                = (35 x 75/100) + (37 x 25/100)

                                                = 35.5

• For simple molecules, such as O₂ and H₂O, the relative molecular mass is calculated by adding the relative atomic masses of the elements involved, giving 32.0 for O₂ and 18.0 for H₂O
• For giant structures, such as Na₂S and SiO₂, the relative formula mass is calculated by adding the relative atomic masses of the elements involved, giving 68.1 for Na₂S and 60.1 for SiO₂

• An orbital is a region that can hold up to two electrons with opposite spins
• Orbitals have different shapes called s, p, d, and f

• S orbitals are spherical in shape and come in sets of one (which can hold up to 2 electrons)

• P orbitals are hour-glass or egg-timer shaped and come in sets of three (which can hold up to 6 electrons)

• D orbitals come in sets of five which can hold up to 10 electrons

• The first energy level (or shell) only contains an s orbital, labelled 1s
• The first shell can hold up to 2 electrons
• The second energy level contains an s orbital (labelled 2s) and three p orbitals  (labelled 2p)
• The second shell can hold up to 8 electrons
• The third energy level contains an s orbital, three p orbitals and five d orbitals
• The third shell can hold up to 18 electrons
• The order in which the orbitals are filled is as follows: 1s 2s 2p 3s 3p 4s 3d 4p
• Note that the 4s fills before the 3d

Evidence that electrons are arranged in shells or energy levels can be obtained by measuring the successive ionisation energies of an element

The first ionisation energy of an element is the energy needed to remove one mole of electrons from one mole of gaseous atoms i.e.

M(g) ® M⁺(g) + e^-

In general, ionisation is easier if the nuclear charge is smaller, the electron is further away from the nucleus and there is more shielding from inner electron shells.

For an element, successive ionisation energies get bigger because the remaining electrons are held more tightly by the unchanged nuclear charge

Jumps occour when going from one shell to another

An element can be assigned to the s, p or d block by working out which type of orbital its outermost electron is in

·         The s block is groups 1 and 2

·         The p block is groups 3 to 8

·         The d block is between the s and p blocks