Acids &redox, Electrons & bonding, Shapes & forces

Question 1

Acids and alkalis

Answer 1

Base- a substance that readily accepts H+ ions from acids
Alkali- a base that dissolves in water, releasing OH- ions in aqeuous solution

Strong acids completely dissociate
Weak acids partially dissociate

Acids are neutralised by bases (metal oxides, metal hydroxides, metal carbonates, ammonia).

Question 2

Preparing standard solution
- what is standard solution

Answer 2

Standard solution is a solution of known concentration.

• Solid is weighed, then dissolved in a beaker using distilled water.
• Transferred to volumetric flask and washings are rinsed into flask.
• Fill with distilled water until bottom of meniscus is touching graduation line.
• Volumetric flask inverted

Question 3

Titration practical

Answer 3

Burette: Contains the solution with the known concentration (titrant). The volume delivered is measured during the titration.

Conical flask: Contains the solution with the unknown concentration. The volume is known and measured beforehand.

• Add a measured volume of one solution (of unknown concentration) to conical flask using pipette.
• Add indicator (e.g. phenolphalein- single universal indicator- colourless to pink).
• Add other solution with unknown concentration to burette and measure intitial burette reading.
• Drop unknown solution into conical flask and swirl, stop when colour change occurs at end point of titration.
• Measure final burette reading- the volume of solution added is called the titre.
• Repeat 3 times and calculate mean. Results should be concordant- 0.1cm^3 apart. Increases accuracy.
• Now you know volume and concentration of burette solution, and the volume of the conical flask solution.
• Use moles to work out unknown concentration.

Question 4

Redox rules
(Sign before number)

Answer 4

Group 1 metals- always +1
Group 2 metals- always +2

O -2 Apart from -1 in peroxides
H +1 Apart from -1 in hydrides
F -1
Cl -1 Apart from when with O or F

Question 5

3 definitions for oxidation

Answer 5

Increase in oxidation number
Loss of electrons
Addition of oxygen

Question 6

What are atomic orbitals? Shape?

How many orbitals in each subshell?
How many orbitals in each shell?

Answer 6

A region around the nucleus that can hold up to 2 electrons with opposite spins.
S orbital- spherical shape
P orbital- dumbell shape

S subshell- 1 orbital
P subshell- 3 orbitals
D subshell- 5 orbitals
F subshell- 7 orbitals

Shell total electrons:
Shell 1: 1s 1 orbital 2 electrons
Shell 2: 2s, 2p 1, 3 orbitals 8 electrons
Shell 3: 3s, 3p, 3d 1,3,5 orbitals 18 electrons
Shell 4: 4s, 4p, 4d, 4f 1,3,5,7 orbitals 32 electrons

Question 7

How do orbitals fill?
Exceptions?

Periodic table arrangement?

Answer 7

• Fill in order of increasing energy
• Orbitals in the same subshell are filled singly first, to prevent repulsion between negatively charged electrons.

4s fills before 3d due to lower energy.
4s also empties before 4d.

Whatever block located in, the highest energy electron is located in that block.

Question 8

Ionic bonding

Answer 8

Electrostatic attraction between oppositely charged ions, in all directions.
Electron transfer from metals and non metals.

Question 9

Properties of ionic compounds

Answer 9

1) h.m.p an h.b.p - strong electrostatic attractions
2) high solubility/ dissolve in polar solvents- polar water molecules break down the lattice and surround the ions.
3) conducts electricity when molten or dissolved:

In solid state:
- Ions are fixed in position in giant ionic lattice
- No mobile charge carriers

Question 10

Quick note- how solubility works

Answer 10

1) Water molecules attract the ions (H+ to -ve, O- to +ve) and the solvent (water) forms intermolecular forces with the ions.
2) This breaks the ionic forces between ions , so ionic lattice breaks down.
2) Water molecules surround the ions, suspending them in solution.

Question 11

Covalent bonding

Dative covalent bonding

Measure of covalent bond strength

Answer 11

The electrostatic attraction between the shared pair of electrons and the nuclei of the bonded atoms
Localised attraction- only occurs between the bonded atoms.

Covalent bond where the shared pair of electrons is supplied by only one of the bonded atoms.
Shown by an arrow.

Average bond enthalpy is the measure of covalent bond strength.

Question 12

Covalent and ionic difference

Answer 12

In covalent the electrons are shared.
In ionic the electrons are transferred.

Question 13

Electron pair repulsion theory

Answer 13

• Electron pairs repel eachother
• so arrange themselves as far apart as possible
• lone pair repels stronger than bonded.

bonded/bonded pair < bonded/lone < lone/lone

Lone pairs are closer to central atom and occupy more space.

Question 14

Bonding region shapes, angles and examples.

How to calculate angle with lone pairs

Answer 14

Bonding region:
2- linear 180 CO2
3- trigonal planar 120 BF3
4- tetrahedral 109.5 CH4
6- octahedral 90 SF6

Find starting angle though no. of bonding regions.
Starting angle- (2.5 x no. of lone pairs)

Question 15

1 lone pair, 3 bonding regions (e.g. NH3)
2 lone pairs, 2 bonding regions (e.g. H2O)
1 lone pair, 2 bonding regions

Answer 15

trigonal pyramidal
Non-linear
Bent

Question 16

Exceptions to shapes and bonds.

Answer 16

Expansion of the octet- more than 8 eectrons in outer shell, e.g. sulfur hexafluoride

Less than 8 electros in outer shell e.g. boron trifluoride

Question 17

Electronegativity

How is it measured

How does it change across the periodic table and why?

Using electronegativity to estimate bonding

Answer 17

The ability of an atom in a molecule to attract the shared pair of electrons in a covalent bond.

By Pauling electronegativity values

Across (right) period table, as charge and size of atoms increases.

Electronegativity: Bonding:
0 covalent
0-1.8 polar covalent
>1.8 ionic

Question 18

Polarity occurs when…

Forces of polar molecules.

Answer 18

Occurs when bonded atoms have different electronegativity values.

A polar covalent bond forms a dipole, called a permanent dipole

Question 19

Dipole definition

Answer 19

Seperation of opposite charges.
Occurs in polar molecules.

Question 20

What are intermolecular forces
Types

Answer 20

Weak interactions between dipoles of different molecules

Hydrogen Bonding (type of permanent)
Permanent dipole-dipole interactions
Induced dipole-dipole interactions (London)

Question 21

London (induced) dipole

Answer 21

Occur between induced dipoles on all molecules

Induced dipole:
Moving electrons produce a changing dipole in a molecule. This is an instantaneous dipole.
This induces a dipole in a neighbouring molecule, carried on…
London forces formed between induced dipoles.

Question 22

Strength of induced dipole

Answer 22

Greater no. of electrons, greater strength.

(because mroe elctrons means larger instantaneous and induced dipole, so the stronger the attractive forces)

Question 23

Permanent dipole-dipole interactions

Answer 23

Occurs between permanent dipole on polar molecules.

Question 24

(Simple molecular substance)

Simple molecular lattice

Answer 24

(e.g., CO2, H2, H20, Ne)- units containing definite no. of atoms.

Simple molecules form a regular structure called a simple molecular lattice. Atoms within each molecule are held together by covalent forces, molecules held together by weak intermolecular forces.

Question 25

Properties of simple molecular substances

Answer 25

–> Low m.p and b.p- weak intermolecular forces break, but covalent bonds do not

–> Non-conductors of electricity- no mobile charge carriers

–> Solubility dependent on polarity:

Non-polar simple molecule in non-polar solvent, intermolecular forces form between the solvent and the molecules.
The interactions weaken the intermolecular forces in lattice, so compound dissolves.

Non-polar simple molecule in polar solvent, little interactions between solvent and molecules in lattice.
Intermolecular forces in simple molecular lattice are too strong to be broken.

Question 26

Hydrogen bonding

Properties of water due to hydrogen bonding

Answer 26

Occurs in molecules containing H attached to a electronegative atom, e.g. N,O,F

• Solid (ice) is less dense than liquid (water)-
  The hydrogen bonds hold the water molecules in a
  fixed positions in an open lattice. Water
  molecules are helf further apart than in water. Ice
  floats in water as it is less dense.
• High m.p an b.p-
  Hydrogen bonds are stronger than london forces, so require more enrgy to break