NCEA Level 3 Thermochemistry

Atomic Radius

The size of atoms are calculated from the average distances between neighbouring atoms

Atomic Radius size is affected by:

Atomic number, electron shells, and electrostatic attraction between positive nucleus and electrons

What happens to Atomic Radius across a period

The effective nuclear charge increases.

Explain why the radius of Cl < Na

Both have the same number of electron shells, the same shielding, the same repulsion, and valence electrons in the same energy level. But Cl has a greater nuclear charge/number of protons, therefor is a stronger attraction for the valence electrons, bringing them closer, resulting in a smaller radius.

What is Electronegativity

Is the measure of attraction between a nucleus and a bonded pair of electrons

What atom has the highest electronegativity

Florine is the most negative.

What does electrongetivity attracition depend on?

Nuclear charge, distance from the nucleus, shielding of inner shells

What does going down a period do to Electronegativity

Decreases

It makes the bonding electrons further away, and there are more layers of shielding electrons, therefore decreasing electronegativity

What does going right do to electronegativity

Increases

Going right on the periodic table increases the number of protons, but the sub orbital doen't change. This increases the strength of attraction between the nucleus and bending electons.

Why is the electronegativity of Cl < F

Cl has more protons that F but it also has more electrons and more electron orbitals. Going down a group, the nuclear charge increases but so does the sheilding, so there is less attraction between the nucleus and bonding pair of electrons.

What is First Ionisation Energy

The minimum energy required to remove one electron from each atom in a mole of atoms in the gaseous state.

Why does ionisation energy follow the trend of electronegativity?

If the electron is more attracted to the nucleus, it will require more energy to remove it.

Electronegativity across a period

INCREASES, The greater the nucleus charge, the easier it is to obtain more electrons from other atoms > more p+ to pull with

1st Ionisation energy - across a period

INCREASES, as the nuclear charge is larger, it requires more energy to remove an electron as they are held tighter (and closer) to the nucleus

Atomic Radii - across a period

DECREAES, across a period the energy level numbers stay the same but nuclear charge increases, pulling more at the valence electrons.

Electronegativity - Down a group

DECREASES, The larger the number of energy levels, the less net electrostatic attraction, the further the neighbouring nucleus so the less ability an atoms has to remove the valance electrons of another atom.

1st Ionisation engergy - Down a group

DECREASES; as the energy level numbers increase, the easier it is to remove electrons, as the valence electrons are further from the 'pull' of the protons and less net electrostatic attraction (and greater radius)

Atomic Radii - Down a period

INCREASES: down a group the energy levels become further away from the nucleus; therefore, the valence electrons repel more, less affected by protons, taking up more space and increasing the size.

Ionic bond

Occurs between a metal and a non-metal atom

Metallic bond

Occurs between atoms in a metal

Covalent bond

Occurs between two non-metal atoms

Covalent bonds form when......

valence electrons are shared. They belong to BOTH atoms simultaneously.

Lewis structure, exceptions

Hydrogen - 2 electrons

Beryllium - 4 electrons

Boron - 6 electrons

More than 8 - anything witha D orbital

Method for drawing Lewis Structures

1. Count the number of valence electrons for all the atoms in the compound

2. Place the central atom

and others around it

3. Connect the atoms with a single bond and subtract this from the total number of valence electrons that you have

4. Add remaining electrons, starting with the outside atoms first before adding to the central atom

5. Check all atoms have full valence shells

6. If central atom doesn't have enough valence electrons. Move non-bonding electrons into a bond, forming a double or triple bond

Polar Bonds

Bonds where the electrons aren't shared evenly

Nonpolar covalent bond

Bonding electrons shared equally between two atoms. No charges on atoms.

Polar covalent bond

Bonding electrons shared unequally between two atoms. Partial charges on atoms.

Ionic bond

Complete transfer of one or more valence electrons. Full charges on resulting ions.

A bond is non-polar when....

2 of the same atoms are bonded to each other

A bond is polar when....

2 different atoms are bonded together

What are the three different types of intermolecular forces

1. Temporary dipole-dipole

2. Permanent dipole-dipole

3. Hydrogen bonds

Permanent dipole-dipole attractions

Between molecules that are polar

Hydrogen bonds

Type of permanent dipole attractions that are much stronger due to their difference in electronegativity; between molecules that have a hydrogen atom bonded to N, O, or F

Strength of temporay dipole-dipole attractions increases as....

Electron cloud size increases/ molecular size increases.

* Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce

As the molecules get bigger the boiling point gets....?

Higher, because there are more temporary dipole-dipole reactions therefore stronger overall intermolecular fore therefore more energy required to intermolecular forces.

Butane (C4H10) and 2-methylpropane (C4H10) have the same mass but different boiling points why? And what has a higher boiling point?

Because butane has more points of contact between molecules, therefore there are more temporary dipole-dipole attractions. therefore increasing intermolecular force therefore more energy is required for break the bonds in butane than 2-methylpropane.

Molecules with permanent dipole-dipole attractions may align so there is an ____________ between the opposite charges on neighbouring molecules. (Because one atom is positive and the other negative)

Electrostatic attraction

In a molecule with hydrogen bonds, what in another molecule is the hydrogen attracted to

the lone pair of either O (oxygen), N (nitrogen), or F (Florine)

The more surface area of the molecule equals a higher boiling point why?

The more surfarce area the more temporary/permanent dp-dp bonds therefore it has a higher boiling point.

If a molecule has hydrogen bonds, it therefore has how many different types of bonds?

All three types

Hydrogen

Permanent

Temporary

Entropy is a measure of?

A measurement of disorder

The more disordered something is, the more entropy it has

Why does volume affect entropy

more volume = more space to move --> increased number of positions/distribution (dispersion of matter and energy), therefore more entropy.

Why does temperature affect entropy

Increased temperature = more kinetic energy therefore particals move faster therefore more disorder therefore more entropy

Why does the number of particles affect entropy

More particals means more positions and increased energy therefore more disorder therefore more entropy

In a boiling point question, what are all the possible types of bonds?

Intermolecular forces

- Temporary dp-dp

- Pernamint dp-dp

- Hydrogen bonds

Ionic Bonds (metal and non-metal)

Metallic Bonds (metal and metal)

Covalent Bonds (non-metal and non-metal)

Why would an Ionic bond be much stronger than an intermolecular bond

An ionic bond is much stronger than an intermolecular bond because it involves a complete transfer of electrons, creating oppositely charged ions that are held together by a powerful electrostatic force

In a boiling point question, what steps should one take?

1. Look for H-bonds

2. Look for polar molecules

3. Look at the atoms of molecules (are they metals)

4. Assign intermolecular forces

5. Check for other bonds (ionic, covalent, metallic)

6. Look at the molar mass of the molecule

7. Look at the shape of the molecule (surface area)