3.1.3 Bonding Flashcards
3.1.3.4 What are the five types of structures?
Monatomic Crystals- Ionic, metallic, macromolecular (giant covalent) and molecular
3.1.3.4 What is absolute zero?
The temp at which particles do not vibrate so have no kinetic heat energy. It is -273degrees Celsius which is 0 Kelvin
3.1.3.4 How is Kelvin calculated?
degrees celsius + 273
3.1.3.4 What is the relationship between temp and KE?
Temperature is directly proportional to the mean kinetic energy of the particles
3.1.3.4 What is latent heat?
The energy required to change state w/o any change in temp (flat line on graph)
3.1.3.4 What happens during condensation/solidifying?
Energy from forming forces/bonds is released as heat energy- exothermic. This stops the temp from falling further as the mean KE of the particles remains constant despite the state change
3.1.3.4 What happens during melting/evaporating?
Heat energy is used to break bonds and surroundings will cool down as heat required for the reaction is absorbed- endothermic
3.1.3.4 Why do liquids cool as they evaporate?
Some KE is used to overcome forces between the particles to allow to allow them to escape so the mean KE of the remaining particles is lower and so the temp is lower.
.5 What is a bonding pair?
The two shared electrons in a covalent bond
.5 What is a lone pair?
Two electrons in a pair not involved in bonding
.5 What is the electron pair repulsion theory?
each pair of electrons around an atom will repel all other electron pairs, the pairs of electrons will therefore lake up positions as far apart as possible to minimise repulsion
.5 What are the five basic shapes of molecules and their bond angles?
Linear (180), trigonal planar (120), tetrahedral(109.5), trigonal bipyramidal (120 and 90) and octahedral (90)
.5 Why do lone pairs repel more than bonding pairs and effect on the bonding angle?
Because they are more compact than bonding pairs so repel more, reducing the bond angles by a small extent
.5 How do you work out the shapes (molecules with single bonds only)?
1) Work out outer shell electrons around central atom 2)Work out number of bonding pairs (no. of atoms bonded to central atom)
3) Work out lone pairs- 1/2 diff between 1 and 2
4) Draw
.5 Examples of linear structures
BeCl2
.5 Examples of trigonal planar structures
ICl3, BH3
.5 Example of tetrahedral structures
AlCl-, BeCl4-, CH4
.5 Example of trigonal bipyramidal structure
PF5
.5 Examples of octahedral structure
SF6
.5 Shapes for 3 electron pairs
3 bonding pairs- Trigonal planar (120) 2bp 1lp bent(V-shape) (118degrees)
.1 What is the formula for the compound ion Sulfate
SO4^2-
.1 Formula for compound ion Nitrate
NO^3-
.1 Formula for compound ion Hydroxide
OH-
.1 Formula for compound ion Carbonate
CO3^2-
.1 Formula for compound ion Ammonium
NH^4+
Ionic Bonding/ Structures
Ionic bonding is the result of electrostatic attraction between oppositely charged ions, caused by the exchanging of electrons to form ions.
The attraction extends throughout the compound.
Every positive ion attracts every negative ion and vice versa.
Ionic electron transfer compounds always exist in a structure called a lattice
MP/BP: High- lots of energy required to overcome strong esf btw ions
EC: When dissolved/molten- ions free to move and carry a current
Strength: Hard and brittle- they form a lattice of alternating positive and negative ions, a blow in the direction shown may move the ions and produce contact between ions with like charges which then repel
Metallic bonding/structure
A lattice of positive ions existing in a ‘sea’ of outer electrons. These electrons are delocalised
Greater charge, stronger attr forces as more electrons released into ‘sea’
Larger in size, weaker attr forces due to greater AR
MP/BP: High- strong esf btw +ve metal ions and -vely charged delocalised electrons
EC: When solid and molten- delocalised electrons free to move and carry a current
Strength: Malleable and ductile- ions/layers in lattice can slide over each other but are still held together by delocalised electrons
Solubility: Insoluble in water- Polarity of water molecs not large enough to overcome strong attractive forces (btw positive ions and negative electrons)
