Periodic Trends

Atomic Size

The electron cloud doesn’t have a definite edge. Scientists get around this by measuring more than 1 atom at a time. Summary: it is the volume that an atom takes up

Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

Group trends – Atomic Radius As we go down a group the atoms have more e-, therefore more energy levels and the atoms get bigger H Li Na K Rb

Periodic Trends – Atomic Radius The atomic radius decreases as you go from left to right across a period. Na MgAl SiPS Cl Ar

Explaining this trend When moving across a period additional p+ are in the nucleus and more e - are in the same energy level. The opposite charges in the nucleus and the e- cloud cause the atom to be 'sucked' together a little tighter. Therefore the radius decreases.

Ionization Energy

Ionization Energy (IE) The amount of energy required to completely remove an e- from a gaseous atom. Recall: removing one e - makes a +1 ion. The energy required is called the first ionization energy. X (g) + energy → X + + e -

Second and third ionization energies represents losing a 2 nd and then a 3 rd e- from the same atom. It can be shown as: X + (g) + energy  X 2+ (g) + e- X 2+ (g) + energy  X 3+ (g) + e- More energy required to remove 2 nd e-, and still more energy required to remove 3 rd e-. The closer the e- is to the nucleus, the more difficult it will be to remove.

Group Trends (I.E.) Ionization energy decreases down the group. Ex. Going from Be to Mg, IE decreases because: –Mg outer e - is in the 3s sub-shell rather than the 2s. –This is higher in energy and further from the nucleus. –So the 3s e - is more easily removed, requiring less energy. A similar decrease occurs in every group in the periodic table.

Period Trends (IE) IE generally increases from left to right. Why? The e - are attracted more strongly to the nucleus (smaller radius). It takes more energy to remove one e - from the atom with stronger attraction, therefore IE increases. Ex. From Na to Ar (11 p+ to 18 p+), the attraction of the protons to e - within the same energy level increases.

Why is there a decrease in IE from Mg to Al? Al is 1s 2 2s 2 2p 6 3s 2 3p 1 It has one e- that is in a p sublevel. Mg is 1s 2 2s 2 2p 6 3s 2. Mg - the ‘s’ sublevel is full – this gives it a slight stability advantage and will require more energy to let go of its e -.

Why is there a fall in IE from phosphorus to sulfur? This can be explained in terms of e- pairing. Phosphorus - 1s 2 2s 2 2p 6 3s 2 3p 3 Sulfur - 1s 2 2s 2 2p 6 3s 2 3p 4 As the p sublevel fills up, e- fill up the vacant sub levels and are unpaired. Phosphorus’ configuration is more energetically stable than sulfur’s because there are e- that are unpaired.

When e- are paired, there is some repulsion which lessens their attraction to the nucleus. It becomes easier to remove! Having a half filled sublevel is more stable than a partially filled sublevel. So… sulfur will break the expected trend and want to lose an e- requiring less IE.

Why an exchange in e - ? Noble Gases have full energy levels. Atoms behave in ways to achieve noble gas configuration.

2 nd Ionization Energy The amount of energy required to remove the 2 nd e- from a gaseous atom. For elements that reach a filled or half filled sublevel by removing 2 e- the 2 nd IE is lower than expected. Makes it easier to achieve a full outer shell True for s 2, the alkaline earth metals which form +2 ions.

3 rd IE Using the same logic s 2 p 1 atoms have a low 3 rd IE. Atoms in the aluminum family form +3 ions. 2 nd IE and 3 rd IE are always higher than 1 st IE!!!

Reactivity

Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily e - can be removed and how strongly atoms want to take other atom's e -.

Reactivity - for Metals: Period - reactivity decreases from left to right Group - reactivity increases going down a group

Why? -Elements located toward the left of the periodic table (alkali metals) and near the bottom easily lose their e -, resulting in higher reactivity. - Within the same group, the more e- an atom has, the easier it will give it up. (Ex. Li (3) and Fr (87))

Reactivity -for Non-Metals Period - reactivity increases from left to right (not including the noble gases) Group - reactivity decreases going down the group. (not including the noble gases)

Why? – Atoms are most stable when they have noble gas electron configuration. –Groups closest to the noble gases want to gain an e- to become stable therefore they have a more vigorous exchange of e-.

–Elements within the same group vary significantly in number of e- but contain the same number of valence e-. –The lower energy levels are found closer to the nucleus, having a stronger desire to complete their energy level and will react more violently.

Shielding Electrons on the outside energy level (valence e-) have the inner energy levels blocking the positive force field (nucleus). These inside energy level e- shield (block) the nuclear (pos) force field from the valence e- shielding.

As you go across the row the nuclear charge (positive charge) gets larger because protons are being added to the nucleus. As you go across the row valence e - are added to the valence shell but the valence e - have the same shielding.

The blocking strength (shielding effect) of these inner e- is the same across the period. Further right in a period the valence e - will have a greater attraction to the nucleus because of the greater positive charge.

Shielding becomes less effective across the row; 2e - can shield +3 better than 2e- can shield +10.

A s you move down a group the valence e- are being added to a new energy level further from the nucleus. These new valence e have additional levels of inner shielding e- and are more effectively shielded from the positive charge. Ex. Campfire

Electronegativity

The tendency for an atom to attract e - to itself when it is chemically combined with another element. How fair it shares. Large electronegativity means it has a strong pull on an e - toward itself.

Group Trend The further down a group the farther the e - is away from the nucleus and the more e - an atom has. Going to the bottom of a group, the e-are further away from the nucleus.

This means they are better shielded from the nuclear (+) charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

Period Trend Electronegativity increases from left to right across a period When the nuclear charge increases, so will the attraction that the atom has for e- in its outermost energy level. This means the electronegativity will increase

Electron Affinity The energy change associated with adding an e - to a gaseous atom. Easiest to add to group 17 or 7A. Gets them to full energy level. Energy is often required (+) when adding an e- to metals. Energy is given off (-) when adding an e- to non-metals. EA decreases as we go down a group.

 Electron affinity decreases as we go down a group because the atoms are getting bigger and the valence electrons are not attracted as strongly to the nucleus.

Ionic Size Cations form by losing e- (have a positive charge). Cations are smaller than the atom they come from. Metals form cations.

Ionic size Anions form by gaining electrons. Anions are bigger than the atom they come from. Nonmetals form anions.

Configuration of Ions Ions always have noble gas configuration. Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.

Periodic Trends Across the period nuclear charge increases so the attractive force gets stronger and the atoms get smaller when filling to the same energy level. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

nimations/chang_7e_esp/pem3s3_2.swf

Size of Isoelectronic ions Iso - same Isoelectronic ions have the same # of electrons Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 all have 10 electrons all have the configuration 1s 2 2s 2 2p 6

Size of Isoelectronic ions Positive ions have more protons so they have a smaller atomic radius. The greater the # of protons the stronger the attraction to the same # of electrons. This will cause the atomic radius to be smaller. Al +3 Mg +2 Na + Ne F-F- O -2 N -3

Organize the isoelectronic ions/atoms in order from smallest to largest P -3, Ar, Cl -, K +, Ca 2+, S 2-, Sc 3+, nimations/chang_7e_esp/pem3s3_4.swfhttp:// nimations/chang_7e_esp/pem3s3_4.swf

Atomic size decreases Ionic size decreases Atomic and Ionicsize increases

Ionization energy, electronegativity, electron affinity INCREASE Ionization energy, electronegativity electron affinity DECREASE