Use LEFT and RIGHT arrow keys to navigate between flashcards;
Use UP and DOWN arrow keys to flip the card;
H to show hint;
A reads text to speech;
25 Cards in this Set
- Front
- Back
|
Ionisation |
The formation of ions, when electrons from atoms are transferred |
|
|
Ions |
1. Carry electric charge 2. Group 4 elements tend to not make ions |
|
|
Ionic compound |
1. Formed when oppositely charged ions are electrostatically attracted to each other 2. EN difference of >1.8 |
|
|
Electronegativity |
Ability of an atom to attract electrons in a covalent bond |
|
|
Ionic crystal lattice |
3D, crystalline structure consisting of oppositely charged ions electrostatically bonded to each other throughout the structure. |
|
|
Covalent bonds |
- Formed by atoms sharing electrons -Seeking to gain electrons to achieve stable electron structure of noble gases - electrostatic attraction between a pair of electrons and positively charged nuclei |
|
|
Types of covalent bonds |
1. bonding pairs of electrons 2. lone pairs of electrons 3. dative pairs (both electrons come from one atom) |
|
|
Exceptions to octet rule |
BF3/BeCl2 - limited space around central atom would cause significant repulsion |
|
|
Short bonds are strong bonds |
Multiple bonds - 1. greater number of shared electrons 2. stronger electrostatic attraction 3. bonded nuclei 4. greater pulling power 5. nuclei brought closer together 6. bonds shorter/stronger than single bonds |
|
|
Polar bonds/Dipole |
1. EN differences (more EN exerts greater pulling power on shared electrons - closer to more EN atom) results in unequal sharing of electrons 2. Indicates bond has 2 separated, opposite, electric charges |
|
|
VSEPR |
Valence shell electron pair repulsion Electron pairs found in the valence shells of atoms repel each other, positioning themselves as far away from each other as possible. |
|
|
Principles of drawing molecules |
1. Draw Lewis structures to identify lone pairs 2. Number of CCs --> shape of molecule 3. Lone pairs will repel other CCs more, because not shared between atoms |
|
|
Linear (No lone pairs - 2 charge centres) |
180º (CO2, BeCl2, C2H2) |
|
|
Planar Trigonal (No lone pairs - 3 charge centres) |
120º (BF3, C2H4) |
|
|
Tetrahedral (No lone pairs - 4 charge centres) |
109.5º (CH4, NH4+, NO3-) |
|
|
Bent (Lone pairs - 3 charge centres) |
117º (SO2) |
|
|
Pyramidal (1 lone pair - 4 charge centres) |
107º (NH3) |
|
|
Bent (2 lone pairs - 4 charge centres) |
105º (H2O) |
|
|
Intermolecular forces |
- determine volatility of substance - strong IMF = more energy required to overcome (^MP/BP) 1. Increasing polarity = increasing IMF 2. Increasing Mr = increasing IMF |
|
|
Van der Waals' |
1) non-polar molecules with no permanent dipole 2) electrons exist in a negatively charged cloud around atoms, can be denser around one atom in one instant 3) instantaneous dipole: weak FOA between opposite ends of temporary dipoles in molecules |
|
|
Overcoming van der Waals' |
1. Low MP/BP 2. Relatively little energy required to overcome FOA + separate molecules 3. Strength of VDW increases with molecular mass |
|
|
Dipole-dipole attraction |
1. molecules with permanent dipole (diff in EN) 2. opposite charges on neighbouring molecules attract each other = dipole-dipole attraction 3. Strength of bond varies according to polarity (HCl > HBr > HI) |
|
|
Hydrogen bonding |
1. H + highly EN atom (F/O/N) a) EN atom pulls electron pair away from H b) hydrogen nucleus very little shielding c) forms strong attractive force with EN atoms of neighbouring molecules' lone electron pairs |
|
|
Hydrogen bonding in ice |
1 H2O has 2H - up to 4 H-bonds can be made. Maximum bonding occurs in ice, causing formation of an open tetrahedral structure, with molecules held fixed distance apart. Results in ice being less dense than water. |
|
|
Metallic bonding |
Positively charged metallic ions in a sea of delocalised, mobile electrons. Arranged in orderly rows |
