Chapter 7 Trends in the Periodic Table

atomic radius ____ along a period

decreases

atomic radius ___ down a group

increases

atomic radius decreases

increase in nuclear effective charge

no increase in screening effect

atomic radius increases

new energy level

screening effect

ionisation/electronegativity ___ along group

increases

ionisation/electronegativity ___ down a group

decreases

ionisation/electronegativity increases

increasing effective nuclear charge

decreasing atomic radius

ionisation/electronegativity decreases

increasing atomic radius

screening effect

alkali metals

reactivity increases down a group, increase in atomic radius screening effect

reaction in water

base and hydrogen gas formed, red litmus paper to blue

halogens

reactivity decreases down group due to electronegativity decreasing

Elements in the same period

• Have the same number of shells (energy levels) outside their nucleus

Elements in the same group

• Have the same number of electrons in their outer shell (energy level)

Group 1 (I) – The Alkali Metals

All the members:

• Have a valency of 1

• Highly reactive

• Become more reactive going down the group

• Are not found freely in nature – only found in compounds with other elements

Properties of the alkali metals:

(1) Soft metals

(2) Low melting points and boiling points

(3) Have low densities

Why are the alkali metals highly reactive?

- They do not satisfy the octet rule – they only have one electron in their outer shell

- In order to satisfy the octet rule they readily lose this electron when they react and form a monopositive ion (+1)

Why does the reactivity of the alkali metals increase going down group I

- The atomic radius increases

Reactions of the alkali metals: (1) Reaction with oxygen

• The alkali metals react rapidly with oxygen to form a metal oxide

• For this reason, they are stored in oil/paraffin

metal + oxygen → metal oxide

Reactions of the alkali metals: 2) Reaction with water

- The alkali metals:

i) Float on water

ii) React vigorously with water

iii) Causes effervescence/fizzing with water

iv) Produce a metal hydroxide and hydrogen gas when reacted with water

metal + water → metal hydroxide + hydrogen

Notes: Metal hydroxides are alkaline – will turn pink with phenolphthalein indicator

To test for Hydrogen gas - Hydrogen gas ignites with a ‘pop

Group 2 (II) – The Alkaline Earth Metals

All the members:

• Have a valency of 2

• Reactive - but not as reactive as the alkali metals

• Become more reactive going down the group

Why are the alkaline earth metals highly reactive?

- They do not satisfy the octet rule – they two electrons in their outer shell

- In order to satisfy the octet rule they readily lose this electron when they react and form a dipositive ion (+2)

Why does the reactivity of the alkali metals increase going down group II

- The atomic radius increases

Reactions of the alkaline earth metals: (1) Reaction with oxygen

• The alkali metals react with oxygen to form a metal oxide

• Not as rapid a reaction as the alkali metals with oxygen

metal + oxygen → metal oxide

Reactions of the alkaline earth metals: (2) Reaction with water

• The alkaline earth metals react with water to form a metal hydroxide and hydrogen gas

• Not as rapid or as vigorous a reaction as the alkali metals with water

metal + water → metal hydroxide + hydrogen

Notes: Metal hydroxides are alkaline – will turn pink with phenolphthalein indicator

To test for Hydrogen gas - Hydrogen gas ignites with a ‘pop’

Transition Elements – the d-block elements

• Transition elements are found between groups II and III in the periodic table

• Transition metals form at least one ion which has an incomplete 3d sublevel

Features of transition elements:

(1) They have variable valencies

Example: Cu+ and Cu2+ (valency of 1 or 2)

(2) They form coloured compounds (not white)

Example: CuSO4 (Copper (II) sulfate) - blue (Cu2+)

Cu2SO4 (Copper (I) sulfate) - brick red (Cu)

(3) They are excellent catalysts

Group 7 (VII) or Group 17 – The Halogens

All the members:

• Have a valency of 1

• Have seven electrons in their outer shell

• Become less reactive going down the group

• Form diatomic molecules – composed of two atoms joined by a covalent bond

Why are the halogens reactive?

- They do not satisfy the octet rule – they seven electrons in their outer shell

• In order to satisfy the octet rule – they gain one electron when they react and form a mononegative ion (-1)

Why does the reactivity of the halogens decrease going down group VII

- The atomic radius increases

What intermolecular forces occur between halogen molecules?

• Van der Waalsforces form between halogen molecules due to being non-polar molecules

Comment on the boiling point of the halogens

• Halogen molecules are non-polar - have weak van der Waals forces between their molecules

• Therefore the halogens have low boiling points

• As the molecule’s molecular mass (Mr) increases, the strength of the van der Waals forces increases

- Fluorine (F2) and Chlorine (Cl2) are gases at room temperature

- Bromine (Br2) is a volatile liquid at room temperature

- Iodine (I2) is a volatile solid at room temperature

Group 8 (VIII) or Group 18 – The Noble Gases

All the members:

• Have a valency of 0

• Have eight electrons in their outer shell - this is a stable arrangement of electrons do not gain, lose, or share electrons – already satisfy the octet rule

• Makes them very inert and safe to use

Examples of using noble gases: Using helium gas in hot air balloons

Using neon gas in lights

Define atomic radius/covalent radius

• Atomic radius is half the distance between the nuclei of two atoms of the same element joined by a single covalent bond

State and explain the trend in atomic radii (covalent radii) across a period of the periodic table

State: Atomic radii values decrease going across a period (atoms get smaller)

Explain: The effective nuclear charge increases - more protons in the nucleus means there is an

increasing attraction between the nucleus and the electrons in the outer shell and the atom is “pulled smaller”

State and explain the trend in atomic radii (covalent radii) down a group of the periodic table

State: Atomic radii values increase going down a group (atoms get larger)

Explain: An additional shell of electrons is added on

Why is establishing an atomic (covalent) radius for the noble gases problematic?

• Atomic radius is half the distance between the nuclei of two atoms of the same element joined by a single covalent bond

• The noble gases satisfy the octet rule – they do not gain, lose, or share electrons so do not form bonds with other elements - they are monatomic

Define electronegativity

• Electronegativity is the measure of relative attraction an atom has for a shared pair of electrons in a covalent bond

State and explain the trend in electronegativity values across a period of the periodic table

State: Electronegativity values increase going across a period (atoms get better at ‘tug of war’!)

Explain:

1. The atomic radius decreases - the pair of electrons in the covalent bond become closer to the positive nuclear charge

2. The effective nuclear charge increases - more protons in the nucleus means there is an increasing attraction between the nucleus and the electrons in the covalent bond

State and explain the trend in electronegativity values down a group of the periodic table

State: Electronegativity values decrease going down a group (atoms get worse at ‘tug of war’!)

Explain:

1. The atomic radius increases - the pair of electrons in the covalent bond become further

from the positive nuclear charge

2. No increase in effective nuclear charge due to an extra shell – electrons in covalent bond are more shielded from the nucleus

Why do the noble gases not have electronegativity values?

• Electronegativity is the measure of relative attraction an atom has for a shared pair of electrons in a covalent bond

• The noble gases satisfy the octet rule – they do not gain, lose, or share electrons so do not form bonds with other elements and so have no electronegativity values

Define ionisation energy

• Ionisation energy is the minimum energy required to remove the most loosely bound electron from an atom or ion

Define first ionisation energy

• First ionisation energy is the minimum energy required to remove the most loosely bound electron from a neutral gaseous atom in the ground state

Representing first ionisation energy in the form of an equation

X → X+ + e-

What unit is used to measure ionisation energy?

• Kilojoules per mole (kJ mol-1)

State and explain the trend in first ionisation energy values across a period of the periodic table

State: First ionisation energy values generally increase going across a period (takes more energy to remove the electron)

Explain:

1. The atomic radius decreases - the electron being removed becomes closer to the positive nuclear charge

2. The effective nuclear charge increases - more protons in the nucleus means there is an increasing attraction between the nucleus and the electron being removed

Important: there are exceptions to this general trend of first ionisation energies going across a period

Notice:

1) Going across a period, first ionisation energy values increase

2) In period 2 and period 3 there are two exceptions to this general increase in first ionisation energy

3) Going from one period to another there is a large decrease in first ionisation energy

State and explain the trend in first ionisation energy values down a group of the periodic table

State: First ionisation energy values decrease going down a group (takes less energy to remove the electron)

Explain:

1. The atomic radius increases - the electron being removed becomes further from the positive nuclear charge

2. No increase in effective nuclear charge due to an extra shell - most loosely bound electron is more shielded from the nucleus

Explaining the exceptions to the general trend in first ionisation energy values across a period

Write out the s, p configuration for the element involved and look at where the first electron isbeing removed from

Electrons being removed from full sublevels, and half-full sublevels will require extra energy as theseare high stability configurations

Explaining the large substantial energy decreases going from one period to another

- An element in the 1st period has their first electron removed from the 1st shell

- An element in the 2nd period has their first electron removed from the 2nd shell

- An element in the 3rd period has their first electron removed from the 3rd shell

➢ It takes substantially less energy to remove an electron from a shell further from the nucleus – this explains why there are large decreases in first ionisation energy going from one period to another

Define second ionisation energy

• Second ionisation energy is the minimum energy required to remove the most loosely bound electron from a positive ion

X+ → X2+ + e-

Give two reasons why the second ionisation energy of an element always greater than the first?

1) The second electron is being removed from a positive ion- the effective nuclear charge has increased so there is a greater attraction between the electron being removed and the nucleus

2) The atomic radius has decreased in the positive ion – the second electron being removed is closer to the positive nuclear charge

Note: If the second ionisation energy of an element is significantly/substantially larger than the first, the second electron is being removed from a shell closer to the nucleus

Looking at successive ionisation energy values for one element

- Questions may be given where the full set of ionisation energy values for a particular element are shown. i.e. the energy values are shown for every electron in the element being removed one by one

- When given successive ionisation energy values for one element, write out the s, p configuration for the element involved and look at where each electron is being removed from

How does looking at successive ionisation energy values for an element provide evidence for the existence of energy levels in atoms?

• When attempting to remove an electron from a new energy level, a substantially higher amount of energy is required.

• Looking at the number of substantial increases in ionisation energy will show how many energy levels are occupied by electrons in the atom