Lecture 12

Question 1

bonds

Answer 1

localised interactions that hold a structure together

Question 2

bond direction

Answer 2

dictates the shape that things adopt

Question 3

bond strength

Answer 3

measure of how hard it is to break a bond, energy is needed to break a bond

Question 4

matter

Answer 4

all matter is held together by bonds, changing the bond type changes its properties

Question 5

solid

Answer 5

strong bonds, short range, complex networks in 3D that hold things in place
lots of bonds need to be broken to remove individual components

Question 6

ordered

Answer 6

crystalline

Question 7

disordered

Answer 7

amorphous

Question 8

liquid

Answer 8

less bonds, weaker
bonds are held close together but no directionality or cooperation
molecules and atoms can slide, making bonds with molecules

Question 9

gas

Answer 9

nothing bonds to anything else, free motionp

Question 10

phase changes

Answer 10

the conversion from one phase to another
bonds are broken because energy (heat) is added

Question 11

ionic solids

Answer 11

rigid 3D network
moving the ions can disrupt many long range bonds
hard, high melting, brittle
insulators, don’t conduct electricity

Question 12

covalent networks

Answer 12

rigid 3D and 2D networks
hard, high melting, brittle
insulators

Question 13

metals

Answer 13

metallic bonding
hard (strong bonds)
ductile (atoms slide but maintain bonding)
conduct electricity

Question 14

intramolecular

Answer 14

strong covalent bonds inside molecules

Question 15

intermolecular

Answer 15

van der waals and hydrogen bonding between molecules

Question 16

covalent bonding

Answer 16

doesn’t take energy to disrupt the weaker bonds
soft, low melting
requires high temperature to start decomposing (break the covalent bond)

Question 17

phase diagram

Answer 17

check notes

Question 18

ionic bonding

Answer 18

electrostatic attraction due to opposite charges between electrons
ions interacting with each other
depend on distance, not direction

Question 19

coulomb’s law

Answer 19

stronger for larger charges and smaller ions

Question 20

ionic lattice

Answer 20

non-directional, larger, 3D arrangements of the ions, ionic lattice form
the packing maximises all of the attractive (-ve to +ve) contacts, as it minimises repulsive charges
the net extra bonding outweighs the extra repulsion

Question 21

result of ionic solid

Answer 21

hard, strong, high melting solids

Question 22

ions

Answer 22

large or small
mixtures of either or both
opposite charges is the only thing that matters

Question 23

ions far apart

Answer 23

cations nucleus repels the anions nucleus and attracts its electrons
cations electrons repel the anions electrons and attract its nucleus

Question 24

medium distances

Answer 24

the gap between them is more important (the atoms are tiny, % change or r is small)
the net positive and negative charges move the ions close together

Question 25

ions close together

Answer 25

attraction doesn't go away the nuclei and some electrons are very close together (very strong repulsion) the electrons that are further away from the nucleus (weaker attraction)

Question 26

ionisation

Answer 26

energy required to remove an electron from an atom

Question 27

electron gain enthalpy

Answer 27

energy required to add an electron to an ion

Question 28

electron transfer example

Answer 28

check notes

Question 29

positive energy

Answer 29

unfavourable dangerous reaction that explodes because it gives out so much energy

Question 30

electron transfer

Answer 30

occurs because electrons 'fall' from a higher energy orbital to a lower energy orbital furthered stabilised by ionic bonds, which are very strong anion always needs a cation

Question 31

valence

Answer 31

charge an ion adopts depends on the position on the periodic table and the no of electrons used in bonding

Question 32

valence orbitals

Answer 32

outer shell is the only one involved in ionic bonding others are too high or low in energy these are the valence shell

Question 33

orbital energies

Answer 33

electron transfer is caused by orbital energy difference the electron falls down that is the potential energy difference

Question 34

electronegativity

Answer 34

measure of a degree to which an atom attracts an electron in a bond increases across a period and decreases down a group

Question 35

inverse relationship

Answer 35

as electronegativity increases, orbitals get lower in energy number that describes the average energy of valence orbitals

Question 36

vaan aarkel triangle

Answer 36

large x = low energy, contacted orbitals small x = radially expanded, high energy orbitals large enough electronegativity leads to an ionic bond.