ATOMIC STRUCTURE AND BONDING Flashcards
Interpreting mass spectra for diatomic molecules
EXAMPLE Cl2
-Two isotope peaks (35Cl+ and 37Cl+)
-Peaks for all possible (3) combinations of these isotopes
Molecular ion peak
Highest value peak
Base peak
Tallest peak (most abundant)
Factors that affect the size of ionisation energy
-Size of nuclear charge
-Size of atom and how many shells it has
-Shielding
-Stability of a 1/2 filled subshell
State and explain the trend in size of atoms across a period
Atomic radius decreases as the number of protons increases, the outer shell electrons become more attracted, resulting in a decrease
Effective nuclear charge explained
-Net charge on the nucleus
-The 10 inner electrons cut down on the outer shell electrons effect, so minus 10 from the number of electrons to find the effective nuclear charge
First ionisation energies decrease down a group because
Whilst nuclear charge increases down a group (1) this is more than often offset by the increase in shielding due to increasing number of inner shells (1). Also because atomic radii increase down a group the outer shell electron is further from the nucleus and are less strongly attracted to it :. requires less energy to remove the electron (1)
First ionisation energies increase across a period because
First ionisation energies generally increase because nuclear charge increases, shielding remains almost constant and atomic radius decreases therefore the outer electron is more strongly attracted to nucleus and requires more energy to remove
Successive ionisation energies explained
Successive removal of electrons generates an increasing positively charged cation which results in a decrease in radii and shielding and the outer shell electrons becoming increasingly attracted to the nuclear charge and making subsequent ionisations progressively more endothermic
SPDF notation- exceptions
Cr- 1s2 2s2 2p6 3s2 3p6 3d5 4s1 (prefers being stable in a half-full subshell)
Cu- 1s2 2s2 2p6 3s2 3p6 3d10 4s1
Transition metals SPDF note
For transition metals, 4s fills first for atoms, 4s empties first for forming ions
For sketching first IE across period 3
-Mg above Na but not above Si
-Al below Mg and above Na
-P above S but below Cl
Strength of the electrostatic forces of attraction increases as…
-The charge on the ions increases
-The size of the ions decrease
Ionic bonding properties
Crystalline- regualr lattic of ions and anions creates crystal structure
High melting and boiling point solids- Large amounts of energy required to break the strong electrostatic attractions between oppositely charged ions in a 3D lattice
Doesnt conduct when solid but does when motlen or aqueous- In the solid state the ions are in fixed positions and cannot move and carry charge, when molten or aqueous, the ions become free to move and carry charge
Tend to be soluable in water- Water collides with the crystal structure breaking the electrostatic attraction and forming ion dipole bonds
Dissolution
The process in which the particles of a substance move into the solvent
Charge density- cations
Simple cations are metal atoms which have lost one or more electrons, they have more protons then electrons hence the positive charge. The ion has a smalller radius than its parent atom because the positive charge in nucleus draws the electrons in
Charge density- anions
Anions have more electrons than protons hence the negative charge. The ion has a larger radius than its parent atom because the extra electron is repelled by neighbouring electrons causing the electron cloud to expand
Atomic radius across and down a group
-Increases down a group (The number of shells increases)
-Decreases across a period (Increasing effective nuclear charge, electron entering the same shell with the same degree of shielding by inner electrons)
The strength of the metallic bond increases as…
-The charge on the positive ion increases
-The size of the positive ion decreases
-The number of delocalised electrons per atom increases
Metallic bonding properties
Hardness- Strong electrostatic forces of attraction between lattice of cations and sea of delocalised electrons
High metling point- Large amount of energy needed to break strong electrostatic forces of attraction between lattic of cations and delocalised electrons
Good electrical conductor- Delocaised electrons are free to move when a potential difference is applied and carry charge
Malleability and ductility- Layers of metal ions can slide over one another, and the delocalise electrons flow into new spaces o maintain metallic bonds
The strength of the covalent bond depends on..
The degree of overlap of atomic orbitals. The greater the amount of overlap, the stronger the bond. Small atoms form strong covalent bonds because their nuclei are close to the bonding electrons resulting in a stronger force of attarction
Properties of simple molecular elements
Low melting and boiling points- The covalent bonds WITHIN the molecule are strong, but the FOA between the molecules are weak. The amount of energy requirede to break these weak intermolecular forces is small
Insulators- There are no free ions or delocalised electrons to move and carry charge
Structure of graphite
Layers of carbon, each carbon is covalntly bound to 3 other carbon atoms (trigonal planar), this leaves one electron delocalised between the layers. Grapite is the only covalent element which conducts electricity
Weak van der waals forces (instantaneous dipole-induced dipole) between layers means that layers can slip over the other when pressure is applied. Used in pencils and lubricants
Diamond
Each carbon atom is covalently bonded to 4 other carbon atoms (tetrahedral) in a 3D structure, resulting in diamond having a very high melting point and is extremely hard and has many strong covalent bonds which require a lot of energy to break
