Module 2: Foundations AS

Question 1

What did J.J Thompson’s “Plum-pudding atom” model suggest?

Answer 1

Atoms are made up of negative electrons moving around in a ‘sea’ of positive charge. This means the overall charge is neutral with equal (+) and (-) charges.

Question 2

What is the shape of an S- orbital?

Answer 2

Spherical

Question 3

Why do ionic compounds have high melting and boiling points?

Answer 3

A large amount of energy is needed to break the strong electrostatic bonds that hold the oppositely charged ions together

Question 4

How do you work out the number of moles from a volume of gas?

Answer 4

moles= V(dm3)

24.0

OR

moles = V(cm3)

24000

NB: Only works at standard atmospheric pressure.

Question 5

What is the Modern day interpretation of the atom?

Answer 5

Negatively charged electrons (-) orbit a nucleus which consists of positively charged protons (+) and neutons which have no charge (neutral/0). Electrons orbit the nucleus on shells. The overall charge of the atom is 0 (Neutral) as the charges of each sub atomic particle (protons, neutrons and electrons) balance/cancel eachother out.

Question 6

What is the relative mass of a single electron?

Answer 6

approximately 1/2000

(Compared to a proton or neutron)

Question 7

What are the different isotopes of Carbon?

Answer 7

¹²₆C (approx 99% abundance)

¹³₆C (approx 1%)

¹⁴₆C (v low; radioactive, formed from cosmic rays bombarding atmospheric Nitrogen)

Question 8

What happens with the electronegativity of elements as we move across the periodic table to Group 7?

Answer 8

Increases

(Elements to right of periodic table need fewer electrons to fill outer shell, so attract electrons more strongly)

Question 9

What volume does one mole of gas occupy at room temperature and pressure?

Answer 9

24.0 dm³

Question 10

What does Electronegativity mean?

Answer 10

The measurment of attraction of a bonded atom for the pair of electrons in a covalent bond.

Question 11

What differences and similarities do isotopes of the same element have?

Answer 11

• Different mases
• The same number of protons and electrons
• Different numbers of neutons in the nucleus
• Undergo the same chemical reactions
• Have (slightly) different physical properties

Question 12

What is the difference between hydrated and anhydrous crystals?

Answer 12

Hydrated crystals contain water whereas anhydrous contain no water

eg

CuSO₄.5H₂O_(s) {blue solid} vs CuSO_4(s) {white solid}

Question 13

Define the term: Non-Polar

Answer 13

When the electrons in a bond are evenly distributed between the atoms that make up the bond; Mainly takes place in covalent bonds between identical atoms (E.g. H-H).

Question 14

What is oxidation?

Answer 14

The GAIN of oxygen and the LOSS of electrons.

This means the compound would be able to give electrons to a reducing substance.

Question 15

What is the oxidation number of combined flourine?

Answer 15

-1 always

(eg NaF or ClF, both have F as -1)

Question 16

What is an oxyanion?

Answer 16

Negative ions that contain an element along with oxygen

eg:

Carbonate: CO₃^2-

Manganate(VII): MnO₄^-

Question 17

Define the term: Polar

Answer 17

When the electrons in a bond are not evenly distributed between bonding atoms (due to one atom being more electronegative/greater attraction); Mainly takes place in covalent bonds between two different atoms.

Question 18

What is elecronegativity?

Answer 18

A measure of the attraction of an electron in a covalent bond

Question 19

What happens when a substance is reduced?

Answer 19

Oxygen is LOST and electrons are GAINED.

The reducing substance gains the electrons from the oxidised substance.

Question 20

What are Intermolecular forces?

Answer 20

Forces which occur between molecules. Applies ONLY to simple covalent molecules.

3 Types:

• van der Waals / London / Dispersion / instantaneous dipole-induced dipole
• Permanent dipole-permanent dipole
• Hydrogen bonding

If these forces are small, the substance will be a gas at room temperature and pressure. The greater the forces, the higher the Melting and Boiling Points.

Question 21

How is electronegativity meassured?

Answer 21

Using the Pauling scale.

This meassures the attraction of a bonded atom for the shared pair of electrons.

Question 22

What are the three main types of intermolecular force?

Answer 22

1. Hydrogen Bonding
2. Van Der Waals’ forces
3. Dipole-Dipole attractions

Question 23

What did Dalton’s atomic theory state?

Answer 23

Atoms are tiny particles that make up elements

Atoms connot be divided

All atoms of a given element are the same

Atoms of an element are different to every other element

Question 24

Do all types of bonds have the same relative strength?

Answer 24

NO!!!

Different types of bond all have different relative strengths…

Question 25

What did Rutherford expect to happen in his gold leaf experiment?

Answer 25

The plum pudding atom would hardly deflect the a-paricles

Question 26

State the correct order of each **sub shell** in terms of **electron configuration**.

Answer 26

1s, 2s, 2p, 3s, 3p, **4s**, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p The main thing to notice is that 4s fills before 3d. Year 13: Ions of transition elements form by losing the 4s electrons first!

Question 27

What did Bohr's model help to explain?

Answer 27

Spectral lines seen in emission spectra Energy of electrons at different distances from the nucleus (In other words, the shell model that you first met at GCSE)

Question 28

Describe metallic bonding

Answer 28

A regular lattice of positive ions surrounded by a sea of delocalised electrons. (The delocalised electrons are from the outer shell of the metal).

Question 29

Describe giant covalent structures and bonding

Answer 29

eg, diamond: Each carbon atom forms 4 strong covalent bonds to neighbouring carbon atoms. The creates a strong 3D lattice. Giant covalent substances are normally insulators with high melting and boiling points. The exception is graphite...

Question 30

Explain why graphite does not behave like a typical giant covalent substance

Answer 30

Graphite is made of carbon atoms arranged into layers. Each carbon atom froms 3 covalent bonds to its neighbours, giving layers of hexagonally arranged carbon atoms. The 4th outer shell electron is delocalised between the layers. Thus the layers can slide over each other, and the delocalised electrons can move between the layers to conduct heat and electricity. (A single layer is called graphene).

Question 31

How does a covalent bond form?

Answer 31

Sharing of electrons brought about by overlap of orbitals from 2 different atoms. Typically covalent bonds form with non-metals only.

Question 32

Describe ionic bonding

Answer 32

Ionic bonding occurs between a metal and a non-metal. The metal loses electron(s) to the non-metal. Each becomes an ion. They are held together in a high melting lattice by strong electrostatic attractions between the oppositely charged ions.

Question 33

How do Hydrogen Bonds form?

Answer 33

Hydrogen bonds can only form if there is a hydrogen atom within a given molecule bonded to eiether N, O or F. The large electronegativity of N, O, F induces a partical positive charge on the H atom, such that it is attracted to a lone pair of electrons from a N, O, F on a neighbouring molecule

Question 34

Describe London forces of attraction

Answer 34

Electrons are constantly in motion. This means that they can be distributed unevenly in a molecule or atom. If this molecule or atom is close to another, this can induce a charge in that molecule, and then the partial charges attract each other. Also called Induced dipole-dipole interactions

Question 35

State the ideal gas law

Answer 35

pV = nRT State **_P_**ressure in **_P_**ascals Volu**_m_**e in **_m_**³ Temperature in Kelvin

Question 36

How do you prepare a standard solution?

Answer 36

Question 37

Define salt

Answer 37

A salt froms when H⁺ ions from an acid are replaced by metal or ammonium ions

Question 38

Give oxidation numbers for the following elements in compounds: O H Na, K Mg, Ca F, Cl, Br, I

Answer 38

O is -2 (except in peroxides: **-1**, eg H₂O₂; with Fluorine **+2**, eg F₂O) H is +1 (except in metal hydrides: -1, eg NaH) Na, K and all group 1 metals are +1 (lose outer shell electron) Mg, Ca and all group 2 metals are +2 (lose 2 outer shell electrons) Halogens are -1 (except Cl, Br, I are positive if bonded to O or a more reactive halogen) eg: IBr (I is +1); IO₂^- (I is +3)

Question 39

State bond angles in water, ammonia and methane

Answer 39

Question 40

Describe p orbitals

Answer 40

Dumbell-shaped: There are 3 p-orbitals in a p sub-shell