Periodic TrendsChapter 6

Octet RuleAtoms tend to achieve electron configuration of Noble GasesOctet = EightNoble Gases have eight electrons in their highest energy levelGeneral Equation for Noble Gases is S2P6

IONSIon- is an atom or group of atoms that have a positive or negative chargeA typical atom is electrically neutral because it has an equal amount of protons and electronsPositive and Negative Ions are formed when at atom donates or receives an electronAn Ion with a positive charge is called a CationAn Ion with a negative charge is called an Anion

Effective Nuclear ChargeForce of attraction between an electron and the nucleus depends on the magnitude of the net nuclear charge acting on the electron and the average distance between the nucleus and the electron.Force of attraction increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.

Effective Nuclear Charge Cont.Valence electron in an atom is attracted to the nucleus of the atom and is repelled by the other electrons.Inner electrons (core) partially shield the outer electrons from the attraction of the nucleus

The effective nuclear charge increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.Periodic Properties: Effective Nuclear ChargeMg has a greater effective nuclear charge than Na, and is smaller than Na.

Trends in Atomic SizeAtomic Radius- the distance between the nuclei of two like atoms in a diatomic moleculeRadius is measured in Picometers1pm = 1 picometer = 1 x 10-12 m

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

Group Trends of Atomic SizeAtomic Size generally increases as you move down a group on the periodic tableAs you descend, electrons are added to higher principle energy levels and the nuclear charge increasesThe outermost orbital is also larger as you move down a groupThe shielding of the nucleus by electrons also increases as you move down a group

Group trendsAs we go down a groupEach atom has another energy level,So the atoms get bigger.HLiNaKRb

ShieldingThe electron on the outside energy level has to look through all the other energy levels to see the nucleus

ShieldingThe electron on the outside energy level has to look through all the other energy levels to see the nucleus.A second electron has the same shielding.

Increasing Atomic Size

Periodic TrendsAtomic Size generally decreases as you move from left to right across a periodAs you go across a period, the energy level remainsEach element has one more proton and electron then the precedingThe electrons are added to the same principle energy level

The effect of the increasing nuclear charge on the outermost electrons is to pull them closer to the nucleusAtomic Size therefore decreases

Periodic TrendsAs you go across a period the radius gets smaller.Same energy level.More nuclear charge.Outermost electrons are closer.NaMgAlSiPSClAr

OverallAtomic NumberAtomic Radius (nm)HLiNeAr10NaKKr

3) Would you expect the atomic radius to be larger or smaller for element 37 than for element 36? Give the reason for your answer.

We would expect that the atomic radius for element 37 to be larger than that of element 36. The trend is for the elements at the beginning of the period (Alkali metals, such as Rb) to have bigger atomic radii than the elements at the end (e.g., Kr). Therefore, the atomic radii diminish as one goes to the right in the periodic table.

CONCLUDING QUESTIONS:

Trends in Ionization EnergyWhen an atom gains or loses an electron, it becomes an ionIonization Energy- The energy required to overcome the attraction of the nuclear charge and remove an electron from an atomThe energy required to remove the first outermost electron is called the first ionization energyThe energy required to remove the second outermost electron is called the 2nd ionization energyEct

Ionization EnergyThe second ionization energy is the energy required to remove the second electron.Always greater than first IE.The third IE is the energy required to remove a third electron.Greater than 1st of 2nd IE.

The Noble Gases are at the top showing they dont want to form an IonThe Alkali are at the bottom of the peaks, showing their ease to form an Ion

SymbolFirstSecond ThirdHHeLiBeBCNO F Ne1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

PROCEDURE "A" QUESTIONS

For any period, the first ionization energy is highest for the noble gas (Group VIIIA) and lowest for the alkali metals (Group IA). The ionization energy then increases from the alkali metals to the noble gases in a period.

8) If there is a periodic variation between first ionization energy and atomic numbers of the elements, how would you describe it?

What determines IEThe greater the nuclear charge the greater IE.Distance form nucleus increases IEFilled and half filled orbitals have lower energy, so achieving them is easier, lower IE.Shielding

Group Trends1st Ionization energy generally decreases as you move down a groupThe size of the atoms increases as you descend, so the outermost electron is farther from the nucleusThe outermost electron should be more easily removed and the element should have a lower ionization energy

Periodic TrendsThe 1st ionization energy generally increases as you move from left to right across a periodThe nuclear charge increases and the shielding effect is constant as you move acrossA greater attraction of the nucleus for the electron leads to the increase in ionization energyExceptions at Full and fill orbitals

First Ionization energyAtomic numberHeHe has a greater IE than H.same shielding greater nuclear charge H

First Ionization energyAtomic numberHHeLi has lower IE than Hmore shielding further awayoutweighs greater nuclear charge Li

First Ionization energyAtomic numberHHeBe has higher IE than Lisame shielding greater nuclear charge LiBe

First Ionization energyAtomic numberHHeB has lower IE than Besame shielding greater nuclear chargeBy removing an electron we make s orbital half filled LiBeB

First Ionization energyAtomic numberHHeLiBeBC

First Ionization energyAtomic numberHHeLiBeBCN

First Ionization energyAtomic numberHHeLiBeBCNOBreaks the pattern because removing an electron gets to 1/2 filled p orbital

First Ionization energyAtomic numberHHeLiBeBCNOF

First Ionization energyAtomic numberHHeLiBeBCNOFNeNe has a lower IE than HeBoth are full,Ne has more shieldingGreater distance

First Ionization energyAtomic numberHHeLiBeBCNOFNeNa has a lower IE than LiBoth are s1Na has more shieldingGreater distanceNa

First Ionization energyAtomic number

Driving ForceFull Energy Levels are very low energy.Noble Gases have full orbitals.Atoms behave in ways to achieve noble gas configuration.

2nd Ionization EnergyFor elements that reach a filled or half filled orbital by removing 2 electrons 2nd IE is lower than expected.True for s2 Alkali earth metals form +2 ions.

3rd IEUsing the same logic s2p1 atoms have an low 3rd IE.Atoms in the aluminum family form + 3 ions.2nd IE and 3rd IE are always higher than 1st IE!!!

Electron AffinityThe energy change associated with adding an electron to a gaseous atom.Easiest to add to group 7A.Gets them to full energy level.Increase from left to right atoms become smaller, with greater nuclear charge.Decrease as we go down a group.

Electron AffinityThe greater the attraction between a given atom and an added electron, the more negative the atoms electron affinityThe more negative the E.A., the greater the attraction of the atom for the electronThe trends in E.A. are not very evident.

Difference between I.E. & E.A.Ionization Energy measures the ease with which an atom loses an electronElectron Affinity measures the ease with which an atom gains an electron

Trends in Ionic SizeAtoms of metallic elements have low ionization energies. They form positive ions easilyAtoms of nonmetallic elements readily form negative ions.How does the lose or gain of electrons affect the size of the ion formed?

Group TrendsPositive Ions are always smaller than the neutral atoms from which they form.The loss of outer-shell electrons results in increased attraction by the nucleus for the fewer remaining electronsNegative Ions are always larger than the neutral atoms from which they formThe effective nuclear attraction is less for an increased number of electrons

Group trendsAdding energy levelIons get bigger as you go down.Li+1Na+1K+1Rb+1Cs+1

The Sodium Atom is larger than the Sodium Cation.Why is this true?

The Chlorine Atom is smaller then the Chlorine Anion.Why is this true?

Sodium Cation is smaller than the Sodium AtomChlorine Anion is larger than the Chlorine Atom

Periodic TrendsGoing from left to right across a row, there is a gradual decrease in the size of the positive ions.Beginning with group 5A, the negative ions, which are much larger, gradually decrease in size an you continue to move right.

Periodic TrendsAcross the period nuclear charge increases so they get smaller.Energy level changes between anions and cations.Li+1Be+2B+3C+4N-3O-2F-1

Trends in ElectronegativityElectronegativity- is the tendency for the atoms of the element to attract electrons when chemically combined with atoms of another element.Electronegativities have been calculated for elements and are expressed in arbitrary units on the Pauling electronegativity scaleThe scale is based on a number of factors

Group TrendsElectronegativity generally decreases as you move down a groupThe metallic elements have a low electronegativity meaning they dont want to want attract electrons

Periodic TrendsAs you go across a period from left to right, the electronegativity of representative elements increasesThe non-metallic elements (excluding Noble Gases) have high electronegativitiesThe trends in electronegativities among transitional metals are not so regular

Electronegativity values help predict the type of ionic or covalent bonding that can exist between atoms in compounds

**************************Figure: 07-12

Title: Electron affinity.

Caption: Electron affinities in kJ/mol for the representative elements in the first five periods of the periodic table. The more negative the electron affinity, the greater the attraction of the atom for an electron. An electron affinity > 0 indicates that the negative ion is higher in energy than the separated atom and electron.**