Unit 3 Flashcards
Binary ionic compound
Composed of ions with metal cation and non metal anion
Ionic bonding definition
the force of attraction between oppositely charged species / ions
- cations and anion oppositely charged so attract each other forming electrostatic attractions to form ionic compounds
- This attraction very strong and requires lots of energy to overcome
Polyatomic ions
They are ions that are made up of more than one type of atom
Formulae of 7 polyatomic ions
Ammonium: NH4+
Hydroxide: OH-
Nitrate: NO3-
Hydrogencarbonate: HCO3-
Carbonate: CO32-
Sulfate: SO42-
Phospate: PO43-
Ionic lattice
Evenly distributed crystalline structure, arranged in regular repeating pattern so positive charge cancels negative charge
Lattice enthalphy
Standard enthalpy change that occurs on the formation of 1 mole of gaseous ions from the solid lattice. Enthalpy change positive value due to endothermic reaction, thus requires to break bonds between ions in lattice
Electronegativity
ability of an atom to draw an electron pair towards itself in a covalent bond
Polar bonds
When atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom
- negative and positive charge don’t coincide, meaning electron distribution is asymmetric
- less electronegative atom gets partial charge of delta positive
- More electronegative atom gets partial charge of delta negative
- Extend of a polarity in covalent bond varies depending on how big the difference exists in the electronegativity values
Covalent bonds and shared electrons
Single bonds: 2 electrons shared
Double bond: 4 electrons shared
Triple bond: 6 electrons shared
Bond energy
energy required to break one mole of a particular covalent bond in the gaseous states. Larger the bond, stronger the covalent bond is
Bond length
distance of two covalently bonded atoms, distance from nucleus of one atom to another.
- Greater the forces of attraction, the more the atoms pulled to each other, this decreases bond length and increases strength.
- Triple bonds are shortest and strongest covalent bonds.
Coordinate bonds
Some molecules have lone pair of electrons which is donated to electron deficient atom which has unfilled outer orbital
What is the VSPER theory
When an atom forms a covalent bond with another atom, the electrons in the different bonds and the non-bonding electrons in the outer shell all behave as negatively charged clouds and repel each other. To minimize this repulsion, all the outer shell electrons spread out as far apart in space as possible
VSPER theory rules
All electron pairs and all lone pairs arrange themselves as far apart in space as is possible.
Lone pairs repel more strongly than bonding pairs.
Multiple bonds behave like single bonds
Dipole
The dipole moment is a measure of how polar a bond is, direction of dipole movement shown by sign in which arrow points to partially negatively charged end of dipole.
how to determine whether a molecule with more than two atoms is polar
- Polar when polar bonds are present, and Non-symmetrical shape (no lone pairs, atoms around central atom same element)
- Non polar when non polar bond or polar bond with symmetrical shape so polarity cancels out
Allotropes
Same element, but exist in different structures. Eg: diamond, graphite, graphene
Simple covalent molecule
number of atoms are fixed everytime.
Giant covalent molecule
number of atoms in a structure can change.
Diamond properties
Tetrahedral 109.5, each carbon bonded to 4 other carbons
All single bonded
Very Strong in strength
Very high Melting/Boiling Point
Not conduct electricity
Graphite properties
Carbon bonded to 3 other carbons, Trigonal planar 120
A toms in the same layer are held together by strong covalent bonds whilst layers connected by weak London Dispersion Force
Very high Melting/Boiling Point
Spare electron delocalised and in space between layers thus can conduct electricity
Fullerene properties
Made of 60 carbons
Trigonal planar, 120
The fourth electron is delocalised so the electrons can migrate throughout the structure
Shape of a football
Graphene
Single layer structure
Very strong
Trigonal planar, 120
Conduct electricity due to delocalized electrons
Silicon
silicon atom bonded to 4 other carbon atoms, Tetrahedral 109.5
All single bonded
Very Strong in strength
Very high Melting/Boiling Point
Not conduct electricity
Silicon dioxide
Tetrahedral at Si, 109.5
Bent at O, 107
High Melting/Boiling Point
Not conduct electricity
Intermolecular forces
forces of attraction between these molecules which must be overcome when the substance is melted and boiled
London dispersion forces
Temporary attraction between temporary dipole, induced dipole of molecules.
- Attraction present in all (polar or nonpolar) covalent molecules
- Only type of Intermolecular force present in non-polar molecules
What does the strength of LDF depend on
LDF proportional to mass of molecule
- larger mass leads to stronger LDF
- Larger surface area with same molecular formula leads to stronger LDF
Dipole dipole attractions
Permanent attraction between polar molecules.
Hydrogen bonding
Permanent attraction between polar molecules, that have H-F, H-O, H-N bonds.
Strongest among the inter molecular force.
physical properties in terms of volatilty for covalent substances
IMF weaker than covalent bonds so substances with low MP/BP are volatile
strength of IMF increase with
- Size of molecule
- Increase in polarity
physical properties in terms of conductivity for covalent substances
- Don’t contain any freely moving particles thus unable to conduct electricity in solid or liquid state
- But some polar covalent molecules ionise and conduct electricity
- Some giant covalent structures capable of conducting electricity due to delocalised electrons
physical properties in terms of solubility for covalent substances
- As covalent molecules larger, solubility decreases as polar part only smaller part of the overall structure
- Giant covalent structures don’t dissolve in solvents as energy needed to overcome strong covalent bonds in lattice structure too great
Metallic bonding
Electrostatic attraction between metal cation & delocalized electrons. Positive charges repel each other and keep the neatly arranged lattice in place. Very strong electrostatic forces between the positive metal centers and the ‘sea’ of delocalized electrons
Properties of metallic bonding
- High melting/boiling point, affected by metallic bond strength which is affected by 1) Ion’s charge 2) Ionic Radius 3) No. of delocalised electrons
- Requires freely moving charge particles for electrical conductivity and occurs in solids and liquids
- Malleable due to attractive forces acting in all direction thus layers slide over and bonds reformed, lattice not broken but changed shape
- Strong and hard due to attractive forces between metal ion and delocalised electrons
Uses of metals examples
- Aluminum is used in food cans because it is non-toxic and resistant to corrosion and acidic food stuffs
- Copper is used in electrical wiring because it is a good electrical conductor and malleable / ductile
- Stainless steel is used for cutlery as it is strong and resistant to corrosion
Alloy
- Mixtures of metals, where the metals are mixed together physically but are not chemically combined. Made from metals mixed with non metals.
- Ions of the different metals are spread throughout the lattice and are bound together by the delocalized electrons
- possible to form alloys because of the non-directional nature of the metallic bonds
Why do alloys have different properties to pure metals
- greater strength, hardness or resistance to corrosion or extreme temperatures
- contain atoms of different sizes, which distorts the regular arrangements of cations. Makes it more difficult for the layers to slide over each other, so they are usually much harder than the pure metal
Common alloys and their uses
Steel (Iron, Carbon): Cars as they are very strong
Bronze (Copper, Tin): Medals due to hard and strong resistance to corrosion
