Chemistry Fundamentals

Question 1

Z

Answer 1

Atomic number

Question 1

A

Answer 1

Mass number

Question 2

Relative atomic mass (Ar)

Answer 2

Mass of one atom relative to 1/12th the mass of 12C

Question 3

Principle quantum number (n)

Answer 3

1st quantum number

Electron shell and size (from 1-7)

Question 4

Angular momentum quantum number (l)

Answer 4

2nd quantum number

Subshell value 0-3

Question 5

4 types of subshell

Answer 5

l - letter - name - electron max
0 - s - ‘Sharp’ - 2 electrons max (Group 1&2)
1 - p - ‘Principal’ - 6 electrons max (Group 13-18)
2 - d - ‘Diffuse’ - 10 electrons max (Transition Metals)
3 - f - ‘Fundamental’ - 14 electrons max (REE)

Question 6

Order of filling shells

Answer 6

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

Question 7

Valence

Answer 7

Valence Shell - Outer Shell
Elements that lose/gain a fixed amount of electrons are monovalent, divalent, trivalent, tetravalent or pentavalent
Valence has no sign (+/-)
Heterovalent means varying valence e.g. transition elements

Question 8

Ionisation Energy

Answer 8

Energy needed for an atom to lose an electron,
Decreases as the principal quantum number increases (further down the periodic table) as the electrons get further from the nucleus and are sheilded by sheels in between.
It also increases from left to right on the periodic table as more electrons are in the shell

Question 9

Electronegativity

Answer 9

Capcity of an atom to attract an electron to itself
measured from 0-4 (4 being the largest)
Electronegativity increases from left to right as atoms have almost comlete outer shells they find it easier to attrct new ones, Electronegativity decreases downwards as there is increase sheilding due to more shells lowering the pull
Low electronegativity is electropositive

Question 10

Ionic Bonding

Answer 10

Electron donation between 2 atoms causes one to become negative (anion) and one to becom positive (cation), these having opposite charges are held together by electrostatic attraction.
Happens in atoms with very different electronegativities
Often forms crystal structures
3 rules: Electrically neutral proportions of ions; distances between cations and anions need to be approx equal to bond length. This is the balance between the attraction between the ions and the repusion between their electron clouds; Each Cation such be surrounded by as many anions as possible and vise-versa in 3d, this is called coordination number.

Question 11

Coordination Number

Answer 11

Number of oppositly charged ions surrounding a central ion
Depends on the radius ratio (Cation size/Anion size)
If the radius ratio is low then CN will be low

Question 12

Substitution of elements

Answer 12

Can happen if: Radius difference between cations is less than 10-15%, sites can strech a bit, the greater the size mismatch the harder it is to substitute; The Charge is the same so the crystal has a net zero charge
Coupled substitution can also occur is 2 elements substitute and balance out their different charges
Trace substituion is common with elements that are not abundant enough to form their own minerals e.g. Ni; Compatible trace elements are those that are preferentially substituted into the crystals so they diminish in the melt with the progression of crystallisation. Incompatible trace elements are those that are enriched in the melt as crystallisation progresses as they cannot fit into the crystals.

Question 13

Magnetic Quantum Number (m)

Answer 13

3rd quantum number
Orbital value from -l to l.
Each orbital can hold up to 2 electrons
Orbitals preferentially fill with a single electron

Question 14

Spin Quantum Number (s)

Answer 14

4th quantum number
Each electron in an orbital can have a spin value of either +1/2 or -1/2.
There can only be one of each in each orbital therefore eplaining why there can only be 2 electrons in an orbital.

Question 15

Covalent Bonding

Answer 15

In elements with similar electronegativities, they cannot form ionic bonds as each atom equally tries to pull the other towards himself. Therefore the atoms share the electrons.
Covalent bonds are the strongest bonds.
In single bonds, 2 orbitals with 1 electron in them from 2 different atoms merge into 1 orbital. This orbital is rugby ball shaped and called a Sigma bond (orbital).
In double and triple bonds, Pi orbitals form as p orbitals merge by overlapping on each side

Question 16

Hybrid orbitals

Answer 16

sp3 hybrid orbitals are when the 1 s orbital and 3 p orbitals are mixed so that each bond has 3/4 p character and 1/4 s character makes all 4 bonds have the same shape.
This is why carbon can form symmetrical tetrahedral shapes with 109 degrees between bonds. This is how SiO4, CH4, and Diamond structures form.
Oxygen can also form an sp3 hybrid orbital even though it has 6 electrons in its outer shell all the electron orbitals a hybridized. These are a special type of sp3 bond with 2 lone electron pairs. These have a higher repulsion force making the angle between the O-H bonds smaller than it should be if it were a symmetrical tetrahedron at 104.5 degrees.
There are also sp3 bonds with 1 lone electron pair. This happens for example in nitrogen and the repulsive force of the lone bond leads to inter bond angles of 107 rather than 109 but larger than the 104.5 of water.
sp2 bonds can also form with 1 s orbital and 2 p orbitals in 1 plane, this makes for flat bonds as is typical in graphite.

Question 17

Ionic-Covalent Continuum

Answer 17

Most bonds are somewhere in between
Depending on the difference electronegativity
Fully covalent bonds are single-element bonds e.g. O-O or C-C this is as they have exactly the same electronegativity
Fully ionic bonds don’t exist as all elements have some electronegativity
Rules:
- The larger the anion for any given cation the more covalent the bond.
- The smaller cation is more covalent for any given anion
- For ions of similar size the one with the largest charge is the most covalent
- Transition metals form more covalent bonds than similar group 2 or 3 elements
For silicates it is important to remember that covalent bonds are stronger so the more covalent a bond is the harder it is to break.

Question 18

Metallic Bonding

Question 19

Weak Bonding

Answer 19

In H20, as the Oxygen atom is more electronegative than the H atoms, the O pulls the bonding electrons closer to itself, leading to the Oxygen being slightly negative and the Hydrogens being slightly positive. As H2O however, is a bent molecule (unlike CO2 for instance) the charges don’t quite cancel each other out. These are very weak charges but important. The shape of the H2O molecule with an angle of 104.5 degrees between the bonds rather than symmetrical makes the water molecule (H2O) into a dipole. A Dipole means it has 2 poles, or a positive and a negative side. The negative side will be the side of the oxygen and the positive side will be the side with the hydrogens.