Structure - Topic 2.1 - Chemistry Flashcards
what is an ionic bond
the strong electrostatic attraction between two oppositely charged ions
state 2 factors which affect the strength of an ionic bond
ionic radii and ionic charge
how does a stronger ionic charge affect the strength of an ionic bond
the greater the charge on an ion, the closer the ions are held together, so stronger ionic bond. This means more energy is needed to break the electrostatic forces, so a higher melting point
how does a smaller ionic radii affect the strength of an ionic bond
smaller ions can pack closer together than larger ions, as electrostatic attraction is larger at a small distance. This means ionic compounds with small ions has a higher melting point
why does ionic radius increase as you go down a group
as you go down the group, the atomic number increases, so extra electrons are added to the shells.
what are isoelectronic ions
ions of different atoms with the same number of electrons
why does the ionic radius of a set of isoelectronic ions decrease as the atomic number increases
as the atomic number increases, the number of electrons stays the same, but the number of protons increases. This means that the electrons are attracted to the nucleus more strongly
what are ionic crystals
giant lattices of ions formed when each ion is electrostatically attracted in all directions to ions of the opposite charge
what are the 4 physical properties of ionic compounds
- high melting point
- soluble in water, but not in non-polar solvents
- don’t conduct electricity when solid
- brittle
what is a molecule
formed when 2 or more atoms bond together by covalent bonds
what is a covalent bond
the strong electrostatic attraction between the two positive nuclei and the shared electrons
between which types of molecules does covalent bonding occur
between non-metals
between which types of molecules does ionic bonding occur
between a metal and a non-metal
where does attraction occur in a covalent molecule
the positive nuclei are attracted to the area of electron density between the two nuclei
where does repulsion occur in a covalent molecule
the two positively charged nuclei and repel each other, as do the electrons
what is bond length
the distance between the 2 nuclei or the distance where the attractive and repulsive forces balance each other out
how does a higher electron density cause a higher bond enthalpy
the higher the electron density, the stronger the attraction between the atoms, so the higher the bond enthalpy
how many electrons are shared in a carbon-carbon triple bond
6 electrons
how many electrons are shared in a carbon-carbon double bond
4 electrons
how many electrons are shared in a carbon-carbon single bond
2 electrons
what is bond enthalpy
how much energy is needed to break a particular chemical bond. The higher the bond enthalpy, the stronger the bond
what is a dative covalent bond
where one atom donates both electrons to a bond
how does type of electron pair affect how much it repels other electron pairs
lone pairs repel more than bonding pairs, meaning bond angles between bonding pairs are reduced, because they are pushed together by lone pair repulsion
what are the 2 factors affecting the shape of the molecule
type of electron pairs surrounding the atom and the number
by how many degrees is the bond angle reduced by when 1 lone pair is added to the molecule
-2.5
state 2 similarities between carbon and silicon
- both form giant covalent structures
- both can form 4 strong, covalent bonds
molecular formula of nitrogen
N2
molecular formula of oxygen
O2
molecular formula of fluorine
F2
molecular formula of Phosphorus
P4
molecular formula of sulfur
S8
molecular formula of chlorine
CL2
describe the structure of diamond
- giant covalent structure
- each carbon atom is bonded to 4 neighbours
- has a tetrahedral arrangement
describe the structure of Silicon (IV) dioxide
- giant covalent structure
- each silicon atom is bonded to 4 neighbours, with oxygen atoms in between
- has a tetrahedral arrangement
state the 5 properties of giant covalent structures
- high melting point
- hard
- good thermal conductors
- insoluble
- can’t conduct electricity
how can graphite conduct electricity
delocalised electrons are free to move along the sheet, thus can carry a charge
what is metallic bonding
when the positive metal ions are electrostatically attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons
why do giant metallic lattices have a high melting point
because of the strong metallic bonding. Also, the more electrons, the stronger the bonding, so the higher the melting point
why are giant metallic lattices malleable and ductile
the layers of metal atoms can slide over each other without breaking the metallic bonds.
why are giant metallic lattices good thermal conductors
the delocalised electrons can pass on the kinetic energy
why are giant metallic lattices good electrical conductors
the delocalised electrons are free to move and can carry a charge
why are giant metallic lattices insoluble
because of the strength of the metallic bonding
what is electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
why does electronegativity increase across a period and up the groups
the more electronegative elements have higher nuclear charges and smaller atomic radii
when is a covalent bond non-polar
when both atoms have similar or identical electronegativities, so the electrons will sit midway between the two nuclei and the bond
why are covalent bonds in homonuclear, diatomic gases (H2, CL2) non-polar
because the atoms have equal electronegativities, so the electrons are equally attracted to both nuclei
when is a covalent bond polar
If the bond is between two electrons with different electronegativities, the bonding electrons will be pulled more towards the more electronegative atom. This causes the electrons to be spear unevenly, causing a dipole
what is a dipole
a difference in charge between the two atoms caused by a shift in electron density in the bond
when is a bond purely covalent
between atoms of a single element, like diatomic gases because the electronegativity difference is zero, so the bonding electrons are arranged evenly
