2.2.2 Bonding and structure Flashcards
define ionic bonding
- ionic bonding is the electrostatic force of attraction between oppositely charged ions in all directions
describe the involvement of outer shell electrons in ionic bonding
- outer shell electrons of metallic atoms are transferred to non-metallic atoms, forming oppositely charged ions
what are some properties of ionic compounds
- solid at room temp
- high mp
- conduct electricity when molten or aqueous but not when solid as ions are free to move and carry charge
- soluble in polar solvents
why do ionic compounds have high mp
- lots of energy required to break the strong electrostatic force of attraction between oppositely charged ions in all directions
what is a giant ionic lattice
regular structure made of the same basic unit repeated over again
> formed because ions are electrostatically attracted in all directions to ions of opposite charged
what 2 factors affect the mp/bp of ionic compounds
- ionic charge: greater charge, stronger electrostatic force, higher mp
- ionic radius: smaller ionic radius, stronger electrostatic force, higher mp
what happens to the ionic radii down the group
- increases down the group
- more electron shells
- ion gets bigger
- lower mp
what happens to ionic radii of isoelectronic ions
- if any ions have same electron configuration
> the ion that is most positive is the smallest
> there are more protons so greater attraction to e- and holds them tighter so higher mp
how are ionic compounds able to dissolve in polar substances
- polar substances e.g. water molecules break down an ionic lattice by surrounding each ion to form a solution
> slight charges within polar substances attract charged ions in giant ionic lattice
why does the greater the ionic charge have reduced solubility
- ionic attraction too strong for water to break down lattice
what is covalent bonding
- the strong electrostatic force of attraction between a shared pair of electrons and the nuclei of bonded atoms
what is the involvement of outer shell electrons in covalent bonding
- outer shell electrons are shared between atoms
a covalent bond can also be described as the overlap…
- overlap of atomic orbitals, each containing one electron to give a shared pair of electrons
what does it mean that the attraction in a covalent bond is localised
- the attraction acts solely between shared pair of electrons and nuclei of the bonded atoms
> this can result in a small unit called a molecule
what is a molecule
- the smallest part of a covalent compound that can exist whilst retaining the chemical properties of the compound
how many covalent bonds do the following form:
Carbon, Nitrogen, Oxygen, Hydrogen
- Carbon - 4
- Nitrogen - 3
- Oxygen - 2
- Hydrogen - 1
what is a dative covalent (coordinate) bond
- a covalent bond where the shared pair of electrons has been supplied by one of the bonding atoms only
> the shared pair was originally a lone pair of electrons on one of the bonded atoms
> represented as an arrow in displayed formula
what is metallic bonding
- the electrostatic force of attraction between metal cations and delocalised electrons
what is the involvement of outer shell electrons in metallic bonding
- outer shell electrons of metallic atoms are lost to form metal cations and a sea of delocalised electrons
what is average bond enthalpy
- a measure of the average energy of a covalent bond
> larger value of average bond enthalpy, stronger the covalent bond
what are the two types of covalent structures
- simple molecular lattice
- giant covalent lattice
describe the bonds + forces within simple molecular structures
- atoms within each molecule are held by strong covalent bonds
- different molecules are held by weak intermolecular forces e.g. london forces
what are the properties of simple molecules
- low mp/bp - weak imf
- non-conductors - no charged particles
- soluble in non-polar solvents
what are the properties of giant covalent stuctures
- high mp/bp - strong covalent bonds
- insoluble in polar + non-polar - covalent bonds too strong to be broken
why do lone pair pf electrons repel more strongly than bonding pairs
- lone pairs of electrons are closer to the central atom and occupy more space than a bonded pair
what is the greatest to least repulsion between bp + lp
- lp/lp
-lp/bp - bp/bp
when drawing a simple molecule / ion what does a solid line represent
- bond in the plane of the paper
when drawing a simple molecule / ion what does a solid wedge represent
- comes out of the plane of the paper (towards you)
when drawing a simple molecule / ion what does a dotted wedge represent
- goes into the plane of the paper (away from you)
what is the shape and bond angle of a molecule with 4 bp + 0 lp
- tetrahedral
- 109.5
what is the shape and bond angle of a molecule with 3 bp + 1lp
- pyramidal
- 107
what is the shape and bond angle of a molecule with 2 bp + 2 lp
- non-linear
- 104.5
what is the shape and bond angle of a molecule with 2 bp + 0 lp
- linear
- 180
what is the shape and bond angle of a molecule with 3 bp + 0 lp
- trigonal planar
- 120
what is the shape and bond angle of a molecule with 6 bp + 0 lp
- octahedral
- 90
what is the electron-pair repulsion theory
- electron pairs surrounding central atom determine the shape of molecule/ion
- electron pairs repel each other so they’re as far apart as possible
- the arrangement of electron pairs minimizes repulsion and holds the bonded toms in a definite shape
what is electronegativity
- the ability of an atom to attract the bonding electrons in a covalent bond
what factors affect electronegativity
- electron shielding
- atomic radius (size of atom)
- nuclear charge
how does nuclear charge affect electronegativity
- as nuclear charge increases, its attraction to electrons on outer shell increases + pulls them closer
> greater electronegativity
how does atomic radius affect electronegativity
- as radius increases, nuclear charge decreases + electron attraction decreases
> smaller atom = greater electronegativity
how does electron shielding affect electronegativity
- more shells so less attraction to electron
> less shells = greater electronegativity
what is the trends in electronegativity down a group
- decreases
> greater electron shielding (inc inner shells)
> atomic radius increase
> nuclear attraction to shared pair of e- decreases
what is the trends in electronegativity across a period
- increases
> nuclear charge increases (more protons)
> atomic radius decreases
> similiar shielding
> nuclear attraction to shared pair of e- increases
in the periodic table how does electronegativity increase
- up and across
what can electronegativity be measured using
- Pauling Scale
what is a non-polar bond (pure covalent bond)
- bonded atoms are the same or have same electronegativity
- electron pair is shared evenly
what is a polar covalent bond
- bonded electron pair is shared unequally
- when bonded atoms are different and have different electronegativities
> the more electronegative atom has greater attraction for bonded pair e-
the higher the difference in electronegativity, the more … a bond
- polar
what is polarity
- uneven sharing of electrons between atoms
how is bond polarity showed
- with partial charges
> delta positive / negative
what is a dipole
- the separation of opposite charges is called a dipole
why is water a polar molecule
- two O-H bond have permanent dipoles
- two dipoles act in diff directions but don’t oppose each other
- O2 = delta negative, H2 = delta positive
why is CO2 a non polar molecule
- two C=O have permanent dipole
- two dipoles act in opposite directions and oppose each other
- over the whole molecule, dipoles cancel out and overall dipole is 0
what are the situations in which dipoles cancel out
- Only when bonds / dipoles are same:
> linear
> trigonal planar
> tetrahedral
> octahedral
what are intermolecular forces
- weak interactions between dipoles of different molecules
intermolecular forces only occur in what
- in simple molecules
What are the 3 different types of Intermolecular forces
- Induced dipole-dipole interactions (London Forces): occurs in all atoms + molecules
- Permanent dipole-dipole interactions: only polar molecules
- hydrogen bonding
describe what london forces are
- weak imf that exist between all molecules (polar + non-polar)
- act between induced dipoles in different molecules
how are london forces formed
- random movement of electrons produces a temporary changing dipole that is constantly changing its position
> known as instantaneous dipole - instantaneous dipole induces dipole on neighbouring molecule
- induced dipole induces further dipoles which then attract each other
> this attraction is the IMF
describe the strength of london forces
- more electrons in molecule means larger the instantaneous and induced dipoles
> therefore greater london forces
> more energy needed to overcome strong interaction so higher bp
what are permanent dipole-dipole interactions
- occurs only in overall polar molecules (that have an overall permanent dipole) in addition to london forces
how are permanent dipole-dipole interactions formed
- interaction between permanent dipoles on different molecules
- molecules that have dipoles but are overall non-polar won’t have these interactions (only london forces)
what is hydrogen bonding
- type of permanent dipole-dipole interaction that occurs for molecules where hydrogen atom is attached to F, O or N
> shape around hydrogen atom is linear
how is hydrogen bonding formed
- acts between a lone pair of electrons on the electronegative atom in one molecule and a hydrogen atom in a different molecule
> forms a hydrogen bond
what is the weakest to strongest of the intermolecular forces
- London forces
- permanent dipole-dipole interactions
- hydrogen bonding
what is broken when simple molecular structures are broken down
- weak IMF holding molecules together break
> covalent bonds between atoms are strong + don’t beak
describe the solubility of non-polar simple molecular structures in non-polar solvents
- tend to dissolve
- imf form between simple molecular compound and solvent molecules, causing simple molecules to break apart from one another and dissolve
describe the solubility of non-polar simple molecular structures in polar solvents
- insoluble
- little interaction between simple molecules and solvent molecules
- imf between polar solvent molecules are too strong to be broken down by interactions with non-polar molecules
describe the properties of simple molecular substances
- low mp/bp - weak imf
- non-electric conductor - no charged particles
how does the density of ice compare with wter
- ice is less dense than water
> hydrogen bonds hold water molecules apart in lattice structure
> water molecules in ice are further apart than in water
> ice floats
why does water have a high mp/bp
- water molecules form hydrogen bonds as well as london forces
> more energy required to break stronger IMF - when ice lattice breaks, the rigid arrangement of hydrogen bonds is broken
> when water boils, the hydrogen bonds break completely
