Trends in the Periodic Table - PowerPoint Presentation

Slide1

Trends in the Periodic Table(Chpt. 7)Slide2

Atomic radius (size) Ionization energy

Electronegativity

The three properties of elements whose changes across the periodic table are to be investigated are:Slide3

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

Slide4

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: Slide5

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)Slide6

Atomic radius

decreases

across the period:

- increasing nuclear charge

- no increase in screening effect

Decrease

Atomic radius

increases

down the group:

- new shell - screening effect

IncreaseSlide7

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 nucleus

No increase in screening effect:

-

same number of shells therefore no increase in screeningSlide8

B) Down a group radius increasesn = 1n = 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 nucleusSlide9

Which is Bigger???

Na or K ?

Na or Mg ?

Al or I ?Slide10

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 minimum

amount of energy required to completely remove the

most loosely bound electron from a neutral gaseous

atom in its ground state

1

st

ionisation energy equation for hydrogen

and sodium

H(g) - e

-

H

+

(g)

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

Slide11

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

-

Na

2+

(g)

Ionisation energy unit – ‘kilojoules per mole’ Table of first ionisation energies given on pg. 80 in the log tablesSlide12

Increase

Decrease

Ionisation Energy

increases

across the period:

- increasing atomic charge

- decreasing atomic radius

Ionisation Energy

decreases

down the group:

- increasing

atomic radius

- screening

effectSlide13

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.Slide14

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. Slide15

Which has a higher 1st ionization energy?

Mg or Ca ?

Al or S ?

Cs or

Ba

?Slide16

Exceptions to the General Trend Across a PeriodIf 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

F

Ne

Atomic Number

First Ionisation Energy

Ionisation Energies for n = 2 periodSlide17

Ionisation Energies for n = 3 period

First Ionisation Energy

Atomic Number

Na

Mg

Al

Si

P

S

Cl

Ar

Slide18

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: 1s

2

,

2s

2 N: 1s2 , 2s2 , 2p3

Mg: 1s

2

, 2s

2

, 2p6 , 3s

2 P: 1s2 , 2s2 , 2p

6 , 3s2 , 3p3 Slide19

*Note: If asked to account fully for trend across 2nd /3rd

period write out electron configurations of Be and N/ Mg and PSlide20

Second and Subsequent Ionisation EnergiesEvidence 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 atomSlide21

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

nucleu

s

b) in a full sublevel

c) has less shieldingSlide22

Investigation of Successive Ionisation energies of Aluminium First 3 ionisation energies increase steadily as

electrons are removed from the 3

rd

shell.

4

th

is a big jump as second shell is entered

5

th

to 11

th

get steadily bigger as successive electrons are removed from the second shell

12th very big as 1st shell is enteredNote: Jumps in ionisation energies are evidence for the existence of the energy levels (shells)Slide23

Trends in ElectronegativityElectronegativity is the relative attraction that an

atom in a molecule has for the shared pair of

electrons in a covalent bond

Concept proposed by

Linus

Pauling

1901-1994Slide24

Increase

Decrease

Electronegativity

increases

across the period:

- increasing nuclear charge

- decreasing atomic radius

Electronegativity

decreases

down the group:

- increasing

atomic radius

- screening effectSlide25

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 electrons

Screening 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.Slide26

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. Slide27

Trends Within GroupsThe 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 propertiesSlide28

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)Slide29

a) Reaction of Alkali Metals with oxygen:All alkali metals react with oxygen to form oxides: Potassium + Oxygen Potassium Oxide

K + O

2

K

2

O

2

½ Slide30

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 + H

2

O

NaOH

+ H

2

½