applied science

How many electrons can an s subshell hold?

2

How many electrons can a p subshell hold?

6

How many electrons can a d subshell hold?

10

Which subshells are available in the first energy level?

s

Which subshells are available in the second energy level?

s and p

Which subshells are available in the third energy level?

s, p and d

What is Hund's rule?

Orbitals must all be singly filled before they can be doubly occupied

Which elements do not fill the 4s subshell before the 3d subshell?

Copper and chromium

Define the term ionic bond

The electrostatic attraction between oppositely charged ions

What is the charge of an ion from group 1?

+1

What is the charge of an ion from group 2?

+2

What is the charge of an ion from group 6?

-2

What is the charge of an ion from group 7?

-1

Explain how atoms of sodium react with atoms of chlorine

Na loses its 2s1 electron gaining a +ve charge.

Cl gains an electron in the 3p subshell gaining a -ve charge.

The opposite charges attract to form NaCl

Why do ionic bonds have such high melting points?

Each +ve ion is surrounded by 6 -ve ions and vice versa.

Strong electrostatic attraction in every direction.

Requires a large amount of energy to break

State two factors that affect the strength of an ionic bond

Size of ion and charge on ion

When can ionic substances conduct electricity?

When molten or in aqueous solution

Describe the properties of ionic compounds

Conduct electricity when molten or aqueous solution

High melting/boiling points

Usually soluble in water

Define the term covalent bond

A shared pair of electrons

Which metals lose electrons from the 4s subshell before the 3d subshell?

Transition metals

Why do metals have such high melting points?

Strong force of attraction between positive ions and delocalised electrons. This requires a large amount of energy to overcome.

State the two factors that affect the strength of metallic bonding

Size of ion

Charge on ion

Explain how the charge on metal ions affects the strength of the metallic bond

The larger the +ve charge the greater the attraction between the nucleus and the delocalised electrons

Explain how the size of the metal ions affects the strength of the metallic bond

The smaller the +ve ion the closer the nucleus is to the delocalised electrons creating a greater attraction

Explain why metals conduct electricity

The delocalised electrons 'carry' charge. Current flows because of this.

Explain why metals conduct heat

Particles are paced tightly so kinetic energy is passed from ion to ion. The delocalised electrons also enable heat to be passed.

Explain why metals are ductile and malleable

The lattice structure allows layers of metal ions to slide over each other without disrupting bonding

Name the 3 forces between molecules

Van der Waals

Permanent dipole-dipole

Hydrogen bonds

Order the 3 forces between molecules in order of strongest to weakest

Hydrogen bonds

Permanent dipole-dipole

Van der Waals

How are Van der Waal's forces formed?

Electrons move to one side, caused temporary dipole. This induces a temporary dipole in neighbouring molecules. Attraction occurs between oppositely charged dipoles

In what molecules do Van der Waal's forces exist?

Non-polar molecules

How are permanent dipole-dipole forces formed?

Permanent dipole in one molecule attracts oppositely charged permanent dipole in neighbouring molecule

In which molecules do permanent dipole-dipole forces exist?

Polar molecules

Which elements must be present for hydrogen bonds to exist?

Hydrogen and either nitrogen, oxygen or fluorine

What is meant by the term displacement?

When a more reactive element takes the place of a less reactive element in a compound

State the equation for determining moles

Moles \= mass ÷ relative atomic mass (molar mass)

Define the term Avogadro's Constant

The number of atoms in a mole of a given substance. Quoted as 6.02x10^23

Define the term relative atomic mass

The average mass of an atom of an element relative to 1/12th the relative atomic mass of Carbon12

Define the term relative molecular mass

The average mass of a molecule relative to 1/12th the relative atomic mass of Carbon12

What does this number represent? 6.02x10^23

The number of particles in a mole. Commonly called Avogadro's Constant

What is the equation for calculating % yield?

% yield \= (actual yield ÷ theoretical yield) x 100

What groups are included in the 's' block of the periodic table?

Groups 1 and 2

What part of the periodic table is known as the 'd' block?

Transition metals

Which groups are in the 'p' block of the periodic table?

3, 4, 5, 6 and 7

What is a group on the periodic table?

A vertical column

What is a period on the periodic table?

A horizontal row

Define the term first ionisation energy

The energy required to remove the outermost electron from one mole of gaseous atoms to produce one mole of gaseous +1 ions

Define the term atomic radius

The distance between the nucleus of an atom and the outermost electron

Define the term electronegativity

A measure of how well an atom attracts a bonding pair of electrons in a covalent bond

Define the term malleability

How easily a material can be hammered into shape

Define the term ductility

How easily a material can be drawn into wires

Describe the trend in atomic radius down any group

Atomic radius increases

Explain the trend in atomic radius down any group

Higher energy levels are filled. The orbitals in higher energy levels are further from the nucleus

Describe the trend in first ionisation energy down groups 1 and 2

First ionisation energy decreases

Explain the trend in first ionisation energy down groups 1 and 2

Increased electron shielding

Greater atomic radius

Smaller attraction to +ve nucleus

SO electron is easier to remove requiring less energy

Describe the trend in melting points down groups 1 and 2

Melting point decreases

Explain the trend in melting points down groups 1 and 2

Strength of metallic bond is weaker due to greater atomic radius decreasing attraction between +ve nucleus and delocalised electrons

Describe the change in state as you go down group 7

The trend is they become more solid (i.e. fluorine is a gas, bromine is a liquid and iodine is a solid)

Describe the change in colour as you go down group 7

They become darker as you go down the group

Describe the trend in electronegativity down group 7

Electronegativity decreases down the group

Explain the trend in electronegativity down group 7

Greater distance between nucleus and bonding electrons

Greater electron shielding

Decreases attraction between nucleus and electron

Describe the trend in melting point down group 7

Melting point increases down the group

Explain the trend in melting point down group 7

Atomic radius increases

Stronger Van der Waal's forces

More energy needed to overcome the intermolecular forces

Describe the trend in atomic radius across a period

Atomic radius decreases

Explain the trend in atomic radius across a period

Greater nuclear charge (more protons)

Same number of electron shells

Same amount of electron shielding

Describe the trend in electronegativity across a period

Electronegativity increases across a period

Explain the trend in electronegativity across a period

Same amount of electron shielding

Greater number of protons

More attraction between nucleus and bonding pair of electrons

Describe the trend in melting point across a period

Melting point increases across the metals and then decreases throughout the non-metals

Explain the trend in melting point across a period

Metallic bonding gets stronger across the period. All other intermolecular forces are weaker than this and therefore easier to break

Why does the group 3 element have a lower first ionisation energy than the group 2 element

Electron taken from p subshell rather than s subshell so is further from the nucleus. Less energy is needed to remove the electron.

Why does the group 6 element have a lower first ionisation energy than the group 5 element

Electron is taken from a paired orbital rather than a singly occupied orbital. Electron repulsion between the pair reduces the energy needed.

Define the term displacement

When a more reactive element takes the place of a less reactive element in a molecule

Describe the trend in reactivity down group 1

They become more reactive down the group

Define the term reduction

Reduction is gain of electrons

Define the term oxidation

Oxidation is loss of electrons

Define the term reducing agent

A reducing agent is something that loses electrons

Define the term oxidising agent

An oxidising agent is something that gains electrons

What oxidation state do group 1 metals have?

+1