Trends in the Periodic Table(Chpt. 7)

1. Atomic radius (size)2. Ionization energy3. Electronegativity

The three properties of elements whose changes across the periodic table are to be investigated are:

1. Trends in Atomic Radii

The atomic radius (covalent radius) of an atom is defined as half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

e.g. in a molecule of hydrogen it is found that the distance between the two nuclei is 0.074 nm.

Therefore the covalent radius of a hydrogen atom is 0.074/2 = 0.037nm

*Note: Noble gases do not form covalent bonds with one another so they have NO atomic radius

The size of an atom depends on the attraction between the positively charged protons and negatively charged electrons in the atom:

- Large attraction – positive protons will pull outer electrons closer to the nucleus - leads to

smaller atomic radius

- Small attraction – electrons will be found further from the nucleus – leads to a larger atomic radius

*Note – understanding the trends in atomic radii values:

Screening (shielding) Effect: means that the inner shell or shells of electrons help to shield the outer electrons from the positive charge in the nucleus

(Please leave space for diagram)

Atomic radius decreases across the period:- increasing nuclear charge- no increase in screening effect

DecreaseAtomic radius increases down the group: - new shell - screening effect

Increase

1. Trends in Atomic radii:

A) Across a period radius decreases 1 2 3 4 5 6 7 8

Increasing effective nuclear charge: - number of protons increases from left to right across a period therefore greater attraction between nucleus and outer electrons – shells drawn closer to nucleusNo increase in screening effect: - same number of shells therefore no increase in screening

B) Down a group radius increases

n = 1

n = 2

n = 3

n = 4

Although there is an increase in the number of protons:

New Shell – additional electrons are going into a new shell which is

further from the nucleus – radius increases

Screening Effect – inner electrons screen outer electrons from the nucleus

Which is Bigger???• Na or K ?• Na or Mg ?• Al or I ?

2. Trends in Ionisation Energy

• Ionisation energy is a term used to describe the tendency of an atom to lose an electron

The First Ionisation Energy of an atom is the minimumamount of energy required to completely remove themost loosely bound electron from a neutral gaseous atom in its ground state

• 1st ionisation energy equation for hydrogen and sodium

H(g) - e- H+(g)

Na(g) - e- Na+(g)

• The second ionisation energy of an element refers to the removal of a second electron from the positive ion formed when the first electron is removed e.g. second ionisation energy of sodium :

Na+(g)

- e- Na2+(g)

• Ionisation energy unit – ‘kilojoules per mole’

• Table of first ionisation energies given on pg. 80 in the log tables

Increase

Decrease

Ionisation Energy increases across the period:- increasing atomic charge- decreasing atomic radius

Ionisation Energy decreases down the group: - increasing atomic radius - screening effect

A) Down a group ionisation energy decreases

Although there is an increase in the number of

protons (nuclear charge):

Increasing atomic radius – radius increases therefore number of shells of electrons increases – outermost electrons are moving further away from attractive force of nucleus and it becomes easier to remove an electron from the outer shell.

Screening Effect – inner electrons screen outer electrons from the positively charged nucleus – becomes easier to remove outermost electrons and ionisation energy values decrease.

B) Across a period ionisation energy increases

Increasing effective nuclear charge – number of protons in nucleus is increasing as move from

left to right across a period. As a result, the attraction between the nucleus and the outer electrons is increasing. More energy is now required to remove an electron from the outermost shell

Decreasing atomic radius – atomic radius decreases from left to right so outer electrons drawn closer to nucleus. Due to increased attraction between electron in outermost shell and nucleus the ionisation energy values increase.

Which has a higher 1st ionization energy?

Mg or Ca ?Al or S ?Cs or Ba ?

Exceptions to the General Trend Across a Period

If plot a graph of ionisation energy Vs atomic number for first 20 elements it is clear that, in any one period, some elements do not follow the smooth increase

Li

Be

B

C

N

O

FNe

Atomic NumberFirs

t Ion

isati

on E

nerg

y

Ionisation Energies for n = 2 period

Ionisation Energies for n = 3 period

Firs

t Ion

isati

on E

nerg

y

Atomic Number

Na

Mg

Al

Si

P

S

Cl

Ar

• In n = 2 period beryllium and nitrogen have higher values than expected

• In the n = 3 period magnesium and phosphorous have higher values than expected

• This irregularity can be explained by the fact that any sublevel that is completely filled (Be, Mg) or exactly half filled (N, P) has extra stability

• Because of this extra stability their ionisation energy values are higher

Be: 1s2 , 2s2 N: 1s2 , 2s2 , 2p3

Mg: 1s2 , 2s2 , 2p6 , 3s2 P: 1s2 , 2s2 , 2p6 , 3s2 , 3p3

*Note: If asked to account fully for trend across 2nd /3rd period write out electron configurations of Be and N/

Mg and P

Second and Subsequent Ionisation Energies

Evidence for the fact that electrons are arranged in shells of different energies is also provided by studying the values of a number of ionisation energies of any one particular element.

Second ionisation higher than the first because removing an electron from:

- an ion so there is more positive charges per electron i.e. greater nuclear charge- closer to nucleus as atomic radius of ion is smaller than corresponding atom

Successive Ionisation Energies- Bigger increase if an electron is removed from a half filled sublevel

- Much bigger if new shell entered because electron being removed is:

a) closer to the nucleusb) in a full sublevelc) has less shielding

Investigation of Successive Ionisation energies of Aluminium

• First 3 ionisation energies increase steadily as electrons are removed from the 3rd shell.

• 4th is a big jump as second shell is entered

• 5th to 11th get steadily bigger as successive electrons are removed from the second shell

• 12th very big as 1st shell is entered

Note: Jumps in ionisation energies are evidence for the existence of the energy levels (shells)

3. Trends in Electronegativity

Electronegativity is the relative attraction that an atom in a molecule has for the shared pair ofelectrons in a covalent bond

Concept proposed byLinus Pauling

1901-1994

Increase

Decrease

Electronegativity increases across the period:- increasing nuclear charge- decreasing atomic radius

Electronegativity decreases down the group: - increasing atomic radius - screening effect

A) Down a group electronegativity decreases

Even though nuclear charge increases down a group:

Increasing atomic radius: - atomic radius

increases – outermost electrons are moving further away from attractive force of nucleus. Therefore smaller attraction between the nucleus and the shared pair of electronsScreening Effect: - inner electrons screen outer electrons from the positively charged nucleus – since it is outermost electrons involved in bonding, the attraction of the nucleus for these electrons decreases going down the group i.e. electronegativity decreases.

B) Across a period electronegativity increases

Increasing effective nuclear charge – number of protons in nucleus is increasing as move from left to right across a period. As a result, the attraction between the nucleus and the outer electrons is increasing. Therefore the electrons involved in bonding are being more strongly attracted to the nucleus i.e. electronegativity increases Decreasing atomic radius – atomic radius

decreases from left to right so outer electrons drawn closer to nucleus. Due to increased attraction between electron in outermost shell and nucleus the

electronegativity values increase.

Trends Within Groups

The chemical properties of elements are largely determined by the number of electrons in the outermost shell:

All elements in group 1 have one electron in outermost shell – all have similar chemical

properties

All elements in group 7 have seven electrons in their outermost shell – all have similar

chemical properties

1. Trends in chemical reactivity of Alkali Metals (Group 1)

• Very reactive elements

• Low ionisation energies and electronegativity values tend to loose electrons to form ionic compounds.

• Reactivity of alkali metals increases down the group (as ionisation energy decreases – more easily outer electron is lost – the more reactive the metal)

a) Reaction of Alkali Metals with oxygen:

All alkali metals react with oxygen to form oxides:

Potassium + Oxygen Potassium Oxide

K + O2 K2 O2 ½

b) Reaction of Alkali Metals with water:

All alkali metals react with water to form the hydroxide of the metal and hydrogen gas is given off:

sodium + water sodium + hydrogen hydroxide

Na + H2O NaOH + H2½